
Class 0/5) 

Book 

Gop$itN? 

CBFHRIGHT DEPOSEE 



A COURSE IN 

INORGANIC CHEMISTRY 

FOR COLLEGES 



BY 
LYMAN C. NEWELL, Ph.D. (Johns Hopkins) 

PROFESSOR OF CHEMISTRY, BOSTON UNIVERSITY, BOSTON, 
MASS., AUTHOR OF c ' EXPERIMENTAL CHEMISTRY" 
" DESCRIPTIVE CHEMISTRY," " GENERAL 
CHEMISTRY," " LABORATORY MAN- 
UAL OF INORGANIC CHEMISTRY 
FOR COLLEGES " 



REVISED 



D. C. HEATH & COMPANY, PUBLISHERS 
BOSTON NEW YORK CHICAGO 



A 



Copyright, 1909, 1916, 
By LYMAN C. NEWELL. 



1g6 



/?/6 



SEP 12 1916 



©CI.A437633 



PREFACE 

This book is intended for college students who devote a 
year to general chemistry. It is primarily a students' book, 
and is not designed to replace the large text-books, which are 
better suited for reference than for class use. The descriptive 
portions of the text include the well-established topics usually 
taught in a year of college chemistry, though considerable 
space is devoted to the application of modern principles to 
chemical industries. The theoretical portions include not 
only the principles whose value was demonstrated long ago, 
but also many recent conceptions, which are fast becoming 
indispensable in interpreting chemical and physical phe- 
nomena. Some of the theoretical topics, which are distributed 
through the text at serviceable points, are the theory of 
electrolytic dissociation, reversible reactions, equilibrium, ca- 
talysis, vapor pressure, electrolysis, and the behavior of dis- 
solved substances. It is hoped that the book as a whole 
will be found adequate for all students whose work is con- 
fined to a year and will likewise serve as a broad foundation 
for those who continue the study of chemistry. 

In the preparation of the manuscript and correction of the 
proof, judicious advice was received from Professor Frank W. 
Durkee, Tufts College, Tufts College, Mass., Assistant Pro- 
fessor William Poster, Princeton College, Princeton, N.J., 
Professor Arthur J\ Hopkins, Amherst College, Amherst, 
Mass., Professor James F. Norris, Simmons College, Boston, 
Mass., Professor Charlotte F. Roberts, Wellesley College, 



iv PREFACE 

Wellesley, Mass., Professor Benjamin W. Van Riper, Wheaton 
College, Wheaton, 111., and others. The author is grateful to 
these teachers for their help, but he assumes entire responsi- 
bility for all statements in the book. 

L. C. N. 
Boston University, Boston, Mass., 
January, 1909. 



PREFACE TO REVISED EDITION 

Opportunity has been taken in this edition to revise the 
portions dealing with the theory of chemistry. Some topics 
have been rewritten and several new ones have been intro- 
duced. This revision and extension include catalysis, osmotic 
pressure, hydrolysis, colloidal solutions, mass action, reversible 
reactions, displacement of equilibrium, solubility product, ad- 
sorption, radioactivity, atomic weights, valence, and molecular 
weights. Many other topics have been improved and extended 
to conform to the advance of science, especially those dealing 
with the applications of chemistry to the arts and industries 
as well as to life itself. Nearly two hundred and fifty new 
problems and exercises have been inserted. Numerical data 
have been corrected and the tables have been revised. These 
extensive modifications, however, have not destroyed the peda- 
gogical features that commended the first edition to such a 

large number of college teachers. 

L. C. N. 
Boston University, Boston, Mass., 
May, 1916. 



CONTENTS 

CHAPTER PAGE 

L Matter, Energy, and Change. • 1 

Properties of Matter — Physical and Chemical Changes 

— Chemical Action — Matter and Energy — Chemical 
Elements — Symbols — Chemical Compounds. 

II. Oxygen — Ozone 15 

Preparation and Properties of Oxygen, and Chemical 
Changes illustrated by Them — Oxides and Oxidation — 
Combustion — Oxygen and Life — Uses — Discovery — 
Ozone. 

III. ^Hydrogen 30 

Preparation and Properties of Hydrogen, and Chemical 
Changes illustrated by Them — Occlusion — Diffusion — 
Oxyhydrogen Blowpipe — Reduction. 

IV. Some Properties of Gases .... 41 

Properties of Gases — Gas Volumes, Temperature, and 
Pressure — Standard Conditions — Law of Charles — 
Absolute Zero — Law of Boyle — Reduction of Gas 
Volumes — Weight of a Liter of Oxygen — Densities of 
Gases. 

V. General Properties of Water ... 50 

Occurrence and Functions in Nature — Industrial Appli- 
cations — Natural, River, and Drinking Waters — Distil- 
lation — Physical Properties — Vapor Tension and 
Vapor Pressure — Chemical Properties — Dissociation by 
Heat — Solvent Power — Solution of Gases — Henry's 
Law — Solutions, Liquids, and Solids — Crystallization 

— Efflorescence — Solution and Vapor Pressure — De- 
liquescence — Thermal Phenomena — Chemical Action. 

v 



vi CONTENTS 

CHAPTER PAGE 

VI. Composition of Water — Hydrogen Dioxide 78 

Composition — Electrolytic Decomposition — Volumetric 
and Gravimetric Composition — Morley's Determination 

— Hydrogen Dioxide. 

VII. Law and Theory . ... . .88 

Law, Theory, Hypothesis — Law of Definite Proportions 

— Law of Multiple Proportions — Atomic Theory — 
Atoms and Molecules — Atomic Weights — Chemical 
Symbols and Formulas — Molecular Weights — Calcula- 
tions — Equations — Four Kinds of Chemical Change — 
Making Equations. 

VIII. Atmosphere — Argon — Nitrogen . . . 115 

Atmosphere — Properties and Ingredients — Properties 
of Nitrogen — Oxygen and Nitrogen in the Atmosphere 

— Volumetric and Gravimetric Composition — Water 
Vapor and Carbon Dioxide in the Atmosphere — Argon, 
Helium, and Related Gases — Air is a Mixture — Liquid 
Air — Liquefaction of Gases — Nitrogen, Preparation, 
Properties, and Relation to Life — Fertilizers. 

IX. Solution — Theory of Electrolytic Disso- 
ciation 132 

General Properties of Solutions — Theory of Electro- 
lytic Dissociation — Ions — Electrolytes and Non-elec- 
trolytes — Osmotic Pressure — Freezing Point and 
Boiling Point — Electrolysis of Solutions — Migration 
of Ions — Chemical Behavior of Electrolytic Solutions — 
Common Ions — Summary. 

X. Acids, Bases, Salts — Neutralization . . 149 

Acids, Properties, Composition, Definition — Bases, 
Properties, Composition, Definition — Salts, Properties, 
Composition — Neutralization — Heat of Neutralization 

— Classification of Salts — Basicity and Acidity — No- 
menclature of Acids, Bases, and Salts — Anhydrides — 
Degree of Dissociation — Salts and Electrolytic Dissoci- 
ation — Hydrolysis. 



CONTENTS vii 

CHAPTER PAGE 

XI. Energy and Chemical Change — Chemical 

Equilibrium ...... 169 

Light and Chemical Change — Heat and Chemical 
Change — Electric Furnace — Measurement of Heat 
Energy — Exothermic and Endothermic Compounds — 
Electricity and Chemical Change — Voltaic and Elec- 
trolytic Cells — Electrotyping and Electroplating — 
Faraday's Law — Catalysis — Equilibrium — Mass Ac- 
tion — Displacement of Equilibrium — Solubility 
Product. 

XII. f Chlorine and its Compounds .... 196 
Chlorine — Chlorine Water — Liquid Chlorine — Bleach- 
ing Powder and Bleaching — Hydrochloric Acid and 
Chlorides — Test for Ionic Chlorine — Chlorine Com- 
pounds. 

XIII. Compounds of Xitrogen — Gay-Lussac's Law 210 

Ammonia — Liquid Ammonia — Ammonium Com- 
pounds — Ammonium Hydroxide — Refrigeration — 
Composition of Ammonia Gas — Nitric Acid and 
Nitrates — Nitrous Acid and Nitrites — Oxides of Ni- 
trogen — Aqua Regia — Gay-Lussac's Law. 

XIV. Atomic and Molecular Weights — Valence 227 

Atomic, Molecular, and Equivalent Weights — Kinetic 
Theory of Gases — Avogadro's Hypothesis — Deter- 
mination of Molecular Weights — Victor Meyer's 
Method — Freezing Point Method — Determination of 
Atomic Weights — Specific Heat — Molecular Formulas 
— Valence. 

XV. Carbon and its Oxides — Carbides . . 263 

Carbon — Diamond — Graphite — Coal — Charcoal — 
Adsorption — Coke — Lampblack — Allotropism — 
Carbon Dioxide — Carbonic Acid — Carbonates — Car- • 
bon Monoxide — Calcium Carbide. 

XVI. Hydrocarbons — Illuminating Gas — Flame 288 

Methane — Ethylene — Acetylene, Illumination , Gen- 
eration — Petroleum — Natural Gas — Coal Gas — 
Water Gas — Illuminating Gases — Elame — Bunsen 
Burner — Oxidizing and Reducing Elame. 



Vlll 



CONTENTS 



CHAPTER 

XVII. Other Carbon Compounds .... 

Organic Compounds — Alcohols — Fermentation — 
Formaldehyde — Ether — Acetic Acid and Acetates 

— Vinegar — Organic Acids — Ethyl Acetate — Fats, 
Glycerine, and Soap — Sugar — Starch — Bread — 
Cellulose — Benzene and Related Compounds — Cy- 
anogen — Proteins. 

XVIII. Sulphur and its Compounds 

Sulphur, Occurrence, Source, Extraction, Purifica- 
tion, Properties — Forms — Transition Point — Hy- 
drogen Sulphide — Sulphides — Sulphur Dioxide — 
Sulphur Trioxide — Sulphurous Acid and Sulphites 

— Sulphuric Acid — Chamber Process — Contact 
Process — Sulphates — Sodium Thiosulphate — Car- 
bon Disulphide. 

XIX. Classification of the Elements — Peri- 
odic Table, Groups, and Law 

XX. Fluorine — Bromine — Iodine 

Fluorine — Hydrofluoric Acid — Bromine — Iodine — 
Hydriodic Acid and Equilibrium — Halogen Family. 

XXI. Boron 

Boron — Boric Acid — Borax — Borax Beads. 

XXII. Silicon — Glass ...... 

Silicon — Silica — Silicic Acids and Silicates — Col- 
loidal Solutions — Silicon Tetrafluoride — Carborun- 
dum — Glass. 



PAGE 

310 



XXIII. 



Arsenic — Antimony — Bis- 



Phosphorus 

muth . . . • . 
Phosphorus — Oxides — Acids and 



332 



355 
362 

374 
378 



391 



Salts — Phos- 
phine — Phosphorus Trichloride — Phosphorus Pen- 
tachloride and its Dissociation — Matches — Kelation 
of Phosphorus to Life — Arsenic — Oxide — Acids 
and Salts — Sulphides — Marsh's Test — Antimony 
and its Compounds — Bismuth and its Compounds 
— Nitrogen Family. 



CONTENTS 



IX 



XXIX. Aluminium 

Occurrence — Metallurgy — Oxide — Gems — Hy- 
droxide — Hydrolysis — Aluminates — Sulphate — 
Alum — Cryolite — Chloride — Clay, Porcelain, and 
Pottery. 



PAGE 

410 



417 



442 



CHAPTER 

XXIV. Metals and Metallurgy .... 

XXV. Sodium — Potassium — Lithium — Ammo- 
nium — Spectrum Analysis . 

Sodium — Chloride — Carbonate — Leblanc and Sol- 
vay Processes — Bicarbonate, Hydroxide, Sulphate, 
Sulphite, Nitrate, Nitrite, and Peroxide — Potas- 
sium — Stassf urt Deposits — Potassium Chloride, 
Nitrate, and Nitrite — Gunpowder — Chlorate, Car- 
bonate, and Hydroxide — Relation to Life — Lith- 
ium — Ammonium Chloride, Sulphate, Sulphide, 
Nitrate, and Carbonate — Spectrum Analysis. 

XXVI. Copper — Silver — Gold .... 

Copper — Metallurgy — Purification by Electrolysis 

— Cuprous Compounds — Cupric Sulphate and Ni- 
trate — Displacement of Metals and Electro-chemical 
Series — Silver — Metallurgy — Silver Plating — 
Silver Nitrate and Halides — Photography — Gold 

— Mining and Metallurgy — Compounds — Copper 
Family. 

XXVII. Calcium — Strontium — Barium — Radium 463 

Calcium — Carbonate and Bicarbonate — Oxide 

— Cement — Hydroxide — Sulphate — Chloride — 
Hardness of Water — Strontium — Barium — Alka- 
line Earth Family — Radium. 

XXVIII. Magnesium — Zinc — Cadmium — Mercury 478 

Magnesium — Oxide, Hydroxide, Sulphate, Chlo- 
ride, and Carbonate — Zinc — Metallurgy — Com- 
pounds — Cadmium — Mercury — Amalgams — 
Compounds — Tests — Zinc Family. 



493 



CONTENTS 

CHAPTER PAGE 

XXX. Tin and Lead — Cerium and Thorium . 505 

Tin — Metallurgy — Transition Point — Alloys — 
Compounds — Oxidation and Reduction — Lead — 
Metallurgy — Drinking Water — Alloys — Oxides, 
Carbonate, Sulphate — Cerium and Thorium. 



XXXI. Manganese 



517 



Preparation, Properties, Uses — Manganese Dioxide 

— Potassium Permanganate — Compounds. 

XXXII. Chromium — Uranium — Radioactivity . 522 

Chromium, Preparation, Properties, Uses — Potas- 
sium Chromate and Dichromate — Chrome Alum 

— Lead Chromate — Compounds — Molybdenum, 
Tungsten, Uranium — Radioactivity. 

XXXIII. Iron — Nickel — Cobalt .... 532 

Occurrence of Iron — Metallurgy — Cast Iron — 
Wrought Iron — Steel, Processes, Properties, Uses — 
Pure Iron — Oxides, Hydroxides, Sulphates, Sul- 
phides, Chlorides, Carbonate, and Cyanides — 
Nickel — Cobalt. 

XXXIV. Platinum and Associated Metals . . 554 



Appendix 



Metric System — Thermometers — Crystallization — 
Vapor Pressure — Atomic Weights. 



557 



A COURSE IN 

INORGANIC CHEMISTRY 

FOR COLLEGES 



INORGANIC CHEMISTRY 



CHAPTER I 

Physical and Chemical Changes — Matter — Energy — Ele- 
ments — Compounds 

Chemistry is one of the natural sciences. By a natural 
science we mean an organized group of knowledge devoted 
primarily to the study of matter and energy. There are 
several of these groups. Some, like geology, deal for the 
most part with the concrete aspects of science, while others, 
like chemistry and physics, are concerned with both abstract 
and concrete phases. The different groups of natural 
sciences are not independent. Indeed, they often overlap 
and have indefinite boundaries. This has always been 
characteristic of chemistry and physics, and recent study 
has drawn them even more closely together. 

The scope of chemistry is rather difficult to outline, espe- 
cially at the outset of its study. Hence it is only possible 
at this stage to present a preliminary and somewhat incom- 
plete conception, leaving to future pages its development, 
interpretation, and application. Chemistry deals with the 
properties of matter, with the changes involved in the trans- 
formations of different kinds of matter, with numerous laws, 
theories, and hypotheses summarizing chemical phenomena, 
and with the manufacture and utilization of a vast number of 
substances indispensable to mankind. 

Properties of Matter. — Different substances are iden- 
tified by characteristics called properties. Thus, some 

1 



2 INORGANIC CHEMISTRY 

substances are solid, others liquid, and still others gaseous, 
though the special physical state of a substance depends 
usually upon its temperature. For example, water, mercury, 
and many other substances, which are liquid at the ordinary 
temperature, become solid if cooled to a low temperature 
and gaseous if heated to a high temperature. Other familiar 
properties are color, odor, taste, relative weight, crystalline 
structure, luster, hardness, melting point, boiling point, and 
solubility. Many properties are exhibited when substances 
are subjected to the action of light, electricity, and par- 
ticularly to a wide range of temperature. But the most 
important properties are doubtless those exhibited when 
different substances act upon each other and thereby produce 
the profound changes characteristic of chemistry. 

Let us consider some properties of the substance sulphur. 
Observation shows it is a solid which has a yellow color, no 
odor, no taste, and a crystalline structure. It is relatively 
heavier than water, because it sinks when placed in a vessel 
of water, but it is insoluble in water. Experiment shows 
that when sulphur is heated, it melts into a pale yellow 
liquid, which turns brown at a comparatively low tempera- 
ture and remains so until the temperature is quite high, 
whereupon it becomes viscous like tar and finally boils, yield- 
ing a yellow smoke which looks much like sulphur. In the 
light sulphur has luster, i.e. it reflects light. It does not 
conduct electricity, for when introduced into the circuit of 
an ordinary electric bell, it prevents the electricity from 
ringing the bell; on the other hand, when rubbed with a 
cloth, sulphur becomes electrified and attracts tiny pieces of 
paper. If a piece of sulphur is heated to a high temperature, 
it takes fire and burns. The flame is blue, and an invisible, 
suffocating product is detected by the odor; if burned long 
enough, all the sulphur is transformed into this gas. Finally, 
if sulphur and a powdered metal, such as iron or copper, are 



PHYSICAL AND CHEMICAL CHANGES 3 

mixed and heated in a test tube, the mixture begins to glow, 
the incandescence often spreading throughout the mass 
even after the test tube has been removed from the flame. 
The product is neither sulphur nor iron but a black substance 
having properties quite different from those of the original 
constituents. Thus, step by step, we have established by 
observation and experiment those properties which serve 
to identify sulphur and to distinguish it from all other sub- 
stances. By a similar though often more elaborate and 
complicated procedure the properties of all substances can 
be discovered, recorded, and classified. 

Changes in Matter and Classification of Properties. — 

Observation and experiment show that substances often 
change or can be changed. These changes are shown by a 
change in properties. Sometimes the change merely in- 
volves a temporary change in properties. Thus, water can 
be changed into steam or ice, and iron can be magnetized or 
melted, but the water and iron are not thereby fundamentally 
altered. Apparently they are changed into new substances, 
but they are not essentially changed, for, although some of 
the properties are new, rigid examination shows that no new 
substances have been formed. The steam, ice, magnetized 
iron, and melted iron are still the substances chemically 
known as water and iron. Indeed, it is only necessary to 
cool the steam and the melted iron and to melt the ice to 
obtain the original substances, which are familiar as water 
and iron. Changes which do not involve the transformation 
of substances into new substances are physical changes. 
In physical changes the change in a substance is often only 
temporary and is due to some change in or departure from 
the conditions which usually prevail, e.g. a change in tem- 
perature, pressure, or electrical conditions, and as soon as 
these new conditions are removed the substance regains its 



4 INORGANIC CHEMISTRY 

familiar properties. Thus, when iodine is heated in a test 
tube, the solid iodine turns into a beautiful violet vapor which 
solidifies in the colder part of the test tube, where it may be 
recognized as iodine by its steel-gray color and crystalline 
form. The iodine has not been transformed into a new sub- 
stance, but is merely changed physically, assuming for the 
time being properties characteristic of iodine at a higher 
temperature than the ordinary. 

Very often, however, substances are changed fundamen- 
tally, their essential nature is affected, and the change is 
manifested by the formation of new substances. The orig- 
inal substance with its properties disappears, often com- 
pletely, and one or more new substances with characteristic 
properties appear. Thus, coal is readily changed into ash 
and an invisible gas which have properties totally unlike 
those of coal. The change is permanent, too, for the ash 
and gas are new substances and do not become coal again as 
soon as they are cold. Changes which involve the trans- 
formation of substances into new substances are chemical 
changes. 

An examination of the properties of sulphur recorded 
above reveals two classes which are more or less distinct. 
Color, odor, taste, relative weight, crystalline structure, 
luster, and electrical behavior are exhibited by unchanged 
or unchanging sulphur. They are associated with physical 
changes, and their manifestation does not involve a dis- 
appearance of the substance nor its transformation into 
another substance. Thus, sulphur is readily recognized as 
sulphur during its immersion in water or its electrification 
by rubbing. Properties exhibited during physical changes 
or by an unchanging substance are called physical properties. 
On the other hand, certain properties — not always readily 
detected — are exhibited only by chemical changes. Their 
manifestation involves a fundamental change in the nature 



PHYSICAL AND CHEMICAL CHANGES 5 

of substances. Thus, sulphur burns in air and unites 
chemically with iron, thereby forming new substances; 
oxygen, as we shall see later, readily undergoes chemical 
transformation, uniting with many substances; further- 
more, many substances are decomposed into other substances 
by heat or electricity. Those properties — not essentially 
physical — manifested by substances when they undergo 
chemical changes are called chemical properties. 

It is clear from the foregoing paragraphs that typical 
physical changes are characterized by the alteration of 
physical properties, while chemical changes are character- 
ized by the formation of one or more new substances. Ex- 
amples of familiar physical changes are changes in physical 
state (illustrated by the formation of ice and vapor from 
water, and vice versa), production of the colors in the sky, 
magnetization of iron, and electrification of glass. Familiar 
chemical changes are the rusting of iron, burning of oil in a 
lamp, digestion of food, souring of milk, and combustion of 
wood. 

It is not always possible to separate physical and chemical 
properties into two sharply defined classes nor to call certain 
changes entirely physical or exclusively chemical. Some 
chemists advocate the adoption of a third class of changes, 
viz. physico-chemical changes, such as fusion {i.e. melting) 
and solution. Thus, when sulphur is boiled or dissolved in 
some liquid, properties are exhibited which are in part physi- 
cal and in part chemical. For example, if boiled sulphur is 
poured while still hot and thick into water, the product is 
a brown, plastic, gummy mass totally unlike the original 
sulphur in appearance though it reverts to ordinary, yellow 
sulphur in time; similarly, powdered sulphur disappears 
{i.e. dissolves) in the liquid carbon disulphide, and is ap- 
parently converted into another substance, though it can be 
readily recovered, often as beautiful yellow crystals, by evap- 



6 INORGANIC CHEMISTRY 

orating the liquid. Many substances besides sulphur when 
fused {i.e. melted) or dissolved exhibit properties which do 
not readily fall into either of the classes mentioned above. 
Such properties are often called physico-chemical properties. 

Chemical Action. — When emphasis is laid upon the chemi- 
cal change, the participating substances or reagents are said 
to undergo chemical action, to interact, or to react. This 
mutual or reciprocal action is called a chemical reaction. 
Usage of these terms is not uniform, but they are used in 
this book in the sense just stated. Thus, when copper is 
put into dilute nitric acid, the evidence of physical change 
in the acid is the rise in temperature, but this is quite sub- 
ordinate to the evidence of chemical change, viz. the gradual 
disappearance of the copper, the liberation of a brown gas, 
and the production of a new substance which dissolves 
readily in the acid and colors it blue. Further experiment 
confirms these observations, for when some of the blue acid 
solution is evaporated, a blue solid is left which dissolves 
in water, gives off the brown gas when heated, and also 
forms a solution which deposits copper upon an iron nail. 
Chemical action has clearly taken place and the whole phe- 
nomenon was a chemical reaction. 

Reactions are often complicated, but extensive study has 
shown that there are four general kinds, — combination, 
decomposition, substitution, and double decomposition. 
These will be illustrated and discussed in appropriate places. 

Matter and Energy. — All changes studied in chemistry 
involve matter and energy. 

It is sufficiently accurate for our present purpose to define 
matter as substance revealed to our senses by properties, the 
fundamental property being weight. In chemical changes 
matter is transformed ; that is, substances are so changed 
that new substances are formed. Such changes, as we have 



PHYSICAL AND CHEMICAL CHANGES 7 

already learned, involve a substitution of new properties for 
old ones. One property of substances, however, is not 
changed, viz. the weight. Careful and extensive study shows 
that the total weight of the substances participating in a 
chemical change is apparently unaltered. This means that 
the sum of the weights of the substances which enter into 
a chemical change equals the sum of the weights of the new 
substances resulting from the chemical change. This feature 
of chemical change is sometimes briefly summed up by saying, 
" matter is indestructible." This rather comprehensive state- 
ment is one form of a law called the Law of the Conservation 
of Matter. It is preferably stated thus: — 

No weight is lost or gained in a chemical change. 

Observation shows that work is done by motion, heat, 
light, and electricity. Thus, at Niagara Falls the water 
turns wheels which are connected with dynamos, and the 
electricity generated thereby is used to operate street cars, 
furnish light, and produce the heat which is used in certain 
chemical industries. The term energy is applied to what- 
ever produces work or can be converted into work. For 
example, motion, heat, light, and electricity are familiar 
forms of energy, and are usually studied in physics. In 
chemistry we study these forms and also another form, called 
chemical energy. Energy, like 'matter, is transformable, 
yet apparently constant in amount. Observation of the 
actual transformations of the different forms of energy is 
an everyday experience. Coal in burning gives up its chemi- 
cal energy in the form of heat, and the heat becomes motion. 
In an electric battery chemical energy is changed into 
electricity, which in turn becomes motion in a bell, or heat 
in a cooking utensil, or light in a metal wire, or even chemi- 
cal energy in an electroplating apparatus. All chemical 
changes involve a transformation of energy. The chemical 
energy is locked up, so to speak, in substances, and when 



8 INORGANIC CHEMISTRY 

chemical changes occur there is a transformation and redis- 
tribution of energy. Thus, the heat which is often needed 
to start chemical changes really becomes transformed, in 
part at least, into chemical energy, while the heat produced 
by chemical changes is the result of the transformation of 
some of the chemical energy into heat. So also the light 
which is essential in photography is transformed into chem- 
ical energy and locked up as such in the chemicals on the 
photographic plate. Food contains chemical energy, and 
when food is digesting, this energy is being transformed into 
heat, which maintains the temperature of the body. But 
all these transformations involve no loss or gain in the total 
amount of energy involved. When a known amount of 
heat is put into a substance and stored up as chemical 
energy, as it were, that heat can be recovered as heat or as 
an equivalent amount of some other form of energy. The 
amount of chemical energy is always altered in chemi- 
cal changes, but the total amount of energy involved 
is unchanged. The characteristics of energy are often 
summarized in a brief statement called the Law of the 
Conservation of Energy, thus : — 

Energy can be transformed without loss, but cannot be (by 
any known means) created or destroyed. 

Chemical Elements and Chemical Compounds. — In chem- 
istry we study almost exclusively two classes of substances, 
viz. elements and compounds. Mixtures, especially the kind 
of mixture known as a solution, receive some attention. 
Chemical compounds consist of chemical elements united 
with one another. There are thousands of compounds, but 
only about eighty elements. 

Chemical Elements. — An extensive examination of many 
different kinds of matter has shown that certain kinds can 
be decomposed at will into two or more substances totally 



PHYSICAL AND CHEMICAL CHANGES 9 

unlike the original matter, but that other kinds are not 
reducible by any means at present under our control. Water, 
for example, can be decomposed into the gases, hydrogen and 
oxygen, which are entirely different from water. But neither 
hydrogen or oxygen can be decomposed by any known process. 
Furthermore, oxygen and hydrogen cannot be transformed 
into each other by any known process. Other substances 
can be added chemically to oxygen or hydrogen, but nothing 
can be taken away from them chemically, nor can anything 
be produced from them without increasing their weight. 
When oxygen or hydrogen undergoes chemical change, there 
is addition, an increase in weight, a transformation into a new 
substance having new properties, and especially a greater 
weight than the original amount of oxygen or hydrogen; 
these chemical changes are never accompanied by loss in 
weight. Furthermore, oxygen and hydrogen have char- 
acteristic properties; some are much alike, while others are 
totally different. Oxygen and hydrogen are elements, and 
all other kinds of matter which have fundamental character- 
istics typified by oxygen and hydrogen are likewise elements. 
The essential properties of the chemical elements will be 
set forth as we proceed with our study of the elements and 
their compounds. It has been customary for many years 
to regard elements as those substances which cannot be 
decomposed into simpler substances or transformed without 
loss of weight. It would follow from this traditional defini- 
tion that the elements are the primary forms of matter, so 
to speak, and are chiefly characterized by stability. This is 
true of most elements, as far as we know. But our inability 
to decompose elements does not necessarily mean that they 
are immutable under all conditions. Indeed, their instability 
is shown by the spontaneous decomposition of certain ele- 
ments, especially radium. Nevertheless, it is sufficient for 
our present purpose to regard the chemical elements as 



10 



INORGANIC CHEMISTRY 



the fundamental materials from which compounds are formed 
and to which compounds are finally reduced. Questions 
concerning transmutation can be postponed until more 
facts are available. 

Each element is designated by a symbol, which is an abbre- 
viation of its name. The following is an alphabetical 



Table of Important Elements and their Symbols 




Name 


Symbol 


Name 


Symbol 


Aluminium 


Al 


Lead . 


Pb 


Antimony 


Sb 


Lithium 


Li 


Arsenic 


As 


Magnesium . . . . . 


Mg 


Barium 


Ba 


Manganese 


Mn 


Bismuth 


Bi 


Mercury 


Hg 


Boron 


B 


Nickel 


Ni 


Bromine 


Br 


Nitrogen 


N 


Cadmium 


Cd 


Oxygen 





Calcium . . . . . . 


Ca 


Phosphorus 


P 


Carbon 


C 


Platinum 


Pt 


Chlorine 


CI 


Potassium 


K 


Chromium 


Cr 


Silicon ...... 


Si 


Cobalt . 


Co 


Silver . 


Ag 


Copper . . . . , . . 


Cu 


Sodium 


Na 


Fluorine 


F 


Strontium . . . . . 


Sr 


Gold 


Au 


Sulphur 


S 


Hydrogen 


H 


Tin 


Sn 


Iodine 


I 


Zinc 


Zn 


Iron 


Fe 







The elements are not uniformly distributed in nature, 
either in abundance or mode of occurrence. Our knowledge 
of the relative abundance of the elements is based on a 
study of the atmosphere, the ocean, and a shell of the 
earth's crust about ten miles deep. The approximate dis- 
tribution of matter in these three portions of the globe is 
seen in the following 



PHYSICAL AND CHEMICAL CHANGES 



11 



Table of the Distribution of Matter 

Atmosphere 03 per cent 

Ocean 7.08 per cent 

Shell of Earth's Crust 92.89 per cent 

The atmosphere contains about 20 per cent of oxygen, 79 
per cent of nitrogen, and 1 per cent of argon. The distribu- 
tion of the elements in the ocean appears in the following 
Table of the Approximate Composition of the Ocean 



Element 



Pee Cent 



Element 



Per Cent 



Oxygen 
Hydrogen . 
Chlorine . 
Sodium 
Magnesium 



85.79 

10.67 

2.07 

1.14 

.14 



Sulphur . . . 
Calcium . . . 
Bromine . . . 
Carbon . . . 
Other Elements 



.09 
.05 
.008 
.002 
traces 



Besides the elements which are combined as saline matter, 
the ocean contains dissolved gases, such as oxygen, nitrogen, 
and carbon dioxide. 

The proportion of the elements in the ten-mile shell of the 
earth's crust is shown in the following 

Table of the Approximate Composition of the Earth's Crust 



Element 


Pee Cent 


Oxygen . . 


47.07 


Silicon . . 


28.06 


Aluminium . 


7.90 


Iron . f . 


4.43 


Calcium . . 


3.44 


Potassium . 


2.45 


Sodium . . 


2.43 


Magnesium 


2.40 


Remainder . 


1.82 



Graphic Proportion 



These elements are chiefly in the combined state. 



12 



INORGANIC CHEMISTRY 



The chemical elements necessary for human beings are 
shown in the following 

Table of the Average Composition of the Human Body 



Element 


Per Cent 


Element 


Per Cent 


Element 


Per Cent 


Oxygen . . . 


65.00 


Phosphorus . 


1.00 


Magnesium . 


0.05 


Carbon . . . 


18.00 


Potassium 


0.35 


Iron . . . 


0.004 


Hydrogen . . 


10.00 


Sulphur . . 


0.25 


Iodine . . . 


trace 


Nitrogen . . 


3.00 


Sodium . . 


0.15 


Fluorine . . 


trace 


Calcium . . 


2.00 


Chlorine . . 


0.15 


Silicon • . . 


trace 



The following is a 

Table of the Less Common Elements and their Symbols 



Name 



Argon . 

Beryllium 

Csesium 

Cerium . 

Columbium 

Erbium 

Europium 

Gadolinium 

Gallium 

Germanium 

Helium . 

Indium . 

Iridium 

Krypton 

Lanthanum 

Molybdenum 

Neodymium 

Neon . . 

Osmium 

Palladium . 

Praseodymium 



Symbol 



A 

Be 

Cs 

Ce 

Cb 

Er 

Eu 

Gd 

Ga 

Ge 

He 

In 

Ir 

Kr 

La 

Mo 

Nd 

Ne 

Os 

Pd 

Pr 



Name 



Radium . 
Rhodium . 
Rubidium 
Ruthenium 
Samarium 
Scandium 
Selenium . 
Tantalum 
Tellurium 
Terbium . 
Thallium . 
Thorium . 
Thulium . 
Titanium . 
Tungsten . 
Uranium . 
Vanadium 
Xenon 
Ytterbium 
Yttrium . 
Zirconium 



Symbol 



Ra 

Rh 

Rb 

Ru 

Sm 

Sc 

Se 

Ta 

Te 

Tb 

Tl 

Th 

Tm 

Ti 

W 

U 

V 

Xe 

Yb 

Yt 

Zr 



PHYSICAL AND CHEMICAL CHANGES 



13 



The symbols of the chemical elements, as may be seen in 
the foregoing tables, are in some instances the first letter 
of the common name of the element. Thus, O is the symbol 
of oxygen, H of hydrogen, N of nitrogen. But since several 
elements have the same initial letter, the symbol of some 
elements contains two letters. Thus, C represents carbon, 
while the symbol of calcium is Ca, of chlorine CI, of chromium 
Cr, and of copper Cu. The symbols of several elements, 
especially the metals so long known, are derived from their 
Latin names. Thus, we have a 

Table of the Chemical Elements with Latin Symbols 



Element 


Latin Name 


Symbol 


E LEMENT 


Latin Name 


Symbol 


Antimony . 


Stibium 


Sb 


Mercury- 


Hydrargyrum 


Hg 


Copper . . 


Cuprum 


Cu 


Potassium . 


Kalium 


K 


Gold . . . 


Aurum 


Au 


Silver . . 


Argentum 


Ag 


Iron . . . 


Ferrum 


Fe 


Sodium . . 


Natrium 


Na 


Lead . . . 


Plumbum 


Pb 


Tin . . . 


Stannum 


Sn 



The symbols of the elements always begin with a capital 
letter, and are not followed by a period. They should be 
learned by actual use. Their full significance w t :11 be ex- 
plained in later chapters. 



Chemical Compounds. — A chemical compound is a sub- 
stance which is composed of two or more chemical elements. 
There is a very large number of chemical compounds. The 
elements which make up a chemical compound are called its 
constituents. Chemical compounds differ fundamentally 
from elements in having three essential characteristics. 
(1) Their constituents are not mixed, but are united chemi- 
cally. That is, the chemical energy originally possessed by 
the elements that make up compounds has so operated that 



14 INORGANIC CHEMISTRY 

the new substance (i.e. the compound) cannot be separated 
into its constituents except by the application of some 
form of energy, such as heat, light, electricity, etc. Thus, 
water is a compound of the elements oxygen and hydrogen, 
which cannot be separated from each other unless water is 
subjected to intense heat or to electricity. (2) In any given 
chemical compound the constituents are always the same 
and are present in a fixed ratio by weight. Thus, sodium 
chloride contains 39.34 per cent of the element sodium and 
60.65 per cent of the element chlorine. (3) The properties 
of a compound differ from those of its constituent elements. 
Thus, the blue solid copper sulphate is a compound of three 
elements, — the red metal copper, the yellow non-metal sul- 
phur, and the colorless gas oxygen. 

Chemical compounds must not be confused with mixtures. 
The components of a mixture may vary in nature and in 
proportion ; they are also not combined chemically but held 
together loosely, and can be separated by some mechanical 
operation, as filtering or sifting, without bringing about any 
chemical change by the application of energy. A mixture, 
too, generally has properties which are similar to, or are an 
average of, those of its components. 

Problems 

1. Solve the problems on the metric system in the Appendix, § 1. 

2. Assuming that the ten-mile shell of the earth's crust weighs 
900 trillion metric tons, calculate the weight of each abundant 
element in it. (See page 11.) 

3. Calculate the weight of each element in a human body weigh- 
ing 70 kilograms. (See page 12.) 

4. Calculate the number of grams in 52 kilograms of typical 
ocean water. 



CHAPTER II 
Oxygen — Ozone 

Oxygen has played an important part in the development 
of chemistry, and is an appropriate element with which to 
begin a systematic study of this science. 

Occurrence. — Oxygen is the most abundant and widely 
distributed of the elements, and occurs both free and com- 
bined. Mixed with nitrogen and a few other gases, it forms 
nearly 21 per cent (by volume) of the atmosphere. Com- 
bined with hydrogen, it constitutes eight ninths (by weight) 
of water; combined with silicon and certain metals, it makes 
up nearly half of the earth's crust; while compounds of 
oxygen with carbon and hydrogen form a large part of 
animal and vegetable matter. Thus, the human body con- 
tains about 65 per cent of oxygen (see page 12), and starch, 
which is a component of all plants, contains 50 per cent. 

Preparation. — Oxygen can be obtained from its compounds 
or from air. It was first prepared by decomposing a com- 
pound of oxygen and mercury called mercuric oxide. When 
heated, this compound decomposes into oxygen and mercury; 
the oxygen escapes as a gas, and the mercury condenses as 
globules or a film on the upper part of the glass vessel in 
which the experiment may be conveniently performed. This 
experiment is historically interesting, because it was first 
performed by Priestley, the discoverer of oxygen. (See 
Discovery of Oxygen, below\) 

15 



16 



INORGANIC CHEMISTRY 



The gas is often prepared by decomposing potassium 
chlorate — a compound of oxygen, chlorine, and potassium. 
If heated to a rather high temperature, the potassium 
chlorate ' passes through a series of changes which result 
finally in the liberation of all the oxygen and the formation 
of a compound of potassium and chlorine called potassium 
chloride. 

Oxygen is most conveniently prepared in the laboratory by 
heating a mixture of potassium chlorate and manganese 




Fig. 1. — Apparatus for preparing and collecting oxygen in the laboratory. 
The oxygen mixture is put in A and heated. The oxygen gas escapes 
through B-C-D into the pneumatic trough (E). The latter is filled with 
water until the shelf is covered and a bottle full of water is then placed 
mouth downward on the shelf. The oxygen bubbles through the hole in 
the shelf up into the bottle and displaces the water. 

dioxide in a glass or metal vessel, and collecting the liberated 
oxygen in a bottle by means of a pneumatic trough. (See 
Fig. 1.) The gas is liberated freely from this mixture at a 
lower temperature than when either compound is heated 
alone. Large quantities of oxygen are prepared for com- 
mercial purposes by heating a mixture of potassium chlorate 
and manganese dioxide in an iron vessel. 



OXYGEN — OZONE 17 

Other commercial processes are used. In Brin's process, 
which is operated largely in England, purified air is forced 
by a pump over barium oxide heated to about 700 degrees 
C., 1 thereby forming barium dioxide. The air supply is 
then cut off, and the pressure in the retorts reduced by 
reversing the pump. This operation changes the barium 
dioxide back again into barium oxide and oxygen. The 
gas is drawn off into a reservoir. The process is then 
repeated. 

Oxygen can be prepared from water. When an electric 
current is passed through water containing sulphuric acid 
or sodium hydroxide, oxygen (and also hydrogen) is liberated. 

Oxygen can also be separated from liquid air. (See Liquid 
Air.) By allowing liquid air to evaporate at ordinary tem- 
perature and pressure, the nitrogen escapes from the liquid 
air more rapidly than the oxygen, leaving finally a liquid 
which is nearly pure oxygen. Unlimited quantities of 
oxygen may probably thus be cheaply prepared from the 
air. 

Oxygen can be prepared from sodium peroxide (or from 
oxone — a special form of this compound) by allowing it 
to interact with water. This is a convenient method for 
preparing a small quantity of the gas. 

Chemical Changes illustrated by the Preparation of Oxy- 
gen. — Brief reference has already been made to the four 
main kinds of chemical change, viz. decomposition, combi- 
nation, substitution, and double decomposition. The first 
and second kinds are illustrated by some of the methods for 
the preparation of oxygen. When mercuric oxide is heated, 

1 C. is the abbreviation of "centigrade," which is the name of the ther- 
mometer used in science. This thermometer registers 100 degrees at the 
boiling point of water, and degrees at the freezing point. (See Appendix, 
§2.) 



18 INORGANIC CHEMISTRY 

it decomposes into oxygen and mercury. Energy in the form 
of heat serves to start and maintain the chemical change, 
and if heat were supplied long enough, all the compound 
would be decomposed into the two elements oxygen and 
mercury. A chemical change whereby a compound with a 
definite set of properties decomposes into two elements each 
with its own set of properties is the simplest illustration of 
the kind of chemical, change called decomposition. This 
fact might be compactly expressed by an equation, thus: — 

Mercuric Oxide = Mercury + Oxygen. 

This equation also emphasizes the fact that the sum of the 
weights of the mercury and oxygen equals the weight of 
the mercuric oxide decomposed — a fact readily proved by 
weighing the three substances. When potassium chlorate 
is heated until all the oxygen is liberated, the chemical 
change is also a decomposition. In this case, however, the 
products are the element oxygen and the compound of 
potassium and chlorine called potassium chloride. The heat 
is not sufficient to produce the complete decomposition of 
potassium chlorate into its three elementary constituents. 
That is, decomposition stops when all the oxygen has been 
liberated. These facts might be expressed thus : — 

Potassium Chlorate = Potassium Chloride + Oxygen. 

Decomposition in general, then, is a chemical change whereby 
a compound is separated chemically into two or more new 
substances. The preparation of oxygen from barium oxide 
and the oxygen of the air illustrates both combination and 
decomposition. In the first stage, when barium oxide is 
heated in a current of air, the compound, barium oxide, 
unites chemically with the element, oxygen, and thereby 
produces a new compound (barium dioxide). A chemical 
change whereby two or more elements or compounds unite 



OXYGEN — OZONE 19 

to form a new compound is called combination. The sec- 
ond stage in the preparation of oxygen by the method under 
discussion is another illustration of decomposition. Thus, 
barium dioxide decomposes into barium oxide and oxygen. 
The two chemical changes might be expressed thus : — 

(1) Barium Oxide + Oxygen = Barium Dioxide. 

(2) Barium Dioxide = Barium Oxide + Oxygen. 

The equations used in this paragraph to summarize the facts 
connected with decomposition and combination are crude 
forms of the chemical equation. Subsequently it will be 
showm that a definite relation exists between the weights of 
the substances used and produced and that this relation can 
be expressed accurately in the form of an equation. At the 
present stage the rather cumbersome equation used above 
is convenient because it emphasizes two facts about chemi- 
cal change; viz. (1) chemical change involves no change in 
the total w r eight of the substances, and (2) chemical change 
results in the formation of new substances. 

Properties. — Pure oxygen gas has no color, odor, or taste. 
It is somewhat soluble in water, one hundred liters of water 
dissolving about three liters of oxygen under ordinary con- 
ditions. The presence of even a small proportion of dissolved 
oxygen is exceedingly important. Fish die in water con- 
taining no dissolved oxygen; and the oxygen absorbed by 
water assists in the decomposition of organic matter into 
harmless gases. Oxygen is slightly heavier than air, its 
density being 1.105 on the air standard. This means that 
under the same conditions of temperature and pressure a 
given volume of oxygen is 1.105 times heavier than an equal 
volume of air. One liter of oxygen weighs 1.429 grams at a 
temperature of 0° C. and under a barometric pressure of 760 
millimeters. Subsequently it will be shown that the volume 



20 INORGANIC CHEMISTRY 

of a gas depends upon the temperature and the pressure^ 
the normal or standard temperature and pressure being 
respectively 0° on the centigrade (C.) thermometer and 760 
millimeters (mm.) on the barometer. Oxygen like other 
gases has three physical states. Liquid oxygen is pale blue. 
It boils at — 182.5 °C. under the normal pressure of the at- 
mosphere; at this temperature its specific gravity is 1.13 {i.e. 
it weighs 1.13 times an equal volume of water). It is mag- 
netic, for when a strong electromagnet is held near it, the 
liquid is attracted by the magnet just as iron filings are by 
an ordinary magnet. Solid oxygen has a whitish color. The 
properties of oxygen enumerated in this paragraph are physi- 
cal properties, because their exhibition does not involve a 
transformation of the oxygen into a new substance. 

The most important chemical property of oxygen is the ease 
with which it combines or interacts with other substances. 
It belongs to the class of active chemical elements, and the 
chemical changes it undergoes are attended by varied and 
interesting physical changes. Oxygen forms compounds 
with all the other single elements except fluorine, bromine, 
argon, and helium (and the other inert gases in the atmos- 
phere). With most of them the union is direct ; i.e. the two 
elements unite chemically to form a compound having oxy- 
gen and the other element as constituents. This direct 
combination is often accompanied by light and heat, though 
the temperature at which combination occurs varies be- 
tween wide limits. At the ordinary temperature oxygen 
unites slowly with most elements, though with phosphorus 
the chemical action is quite rapid, as may be seen by the 
glow and fumes when a phosphorus-tipped match is rubbed, 
especially in the dark. Metals, such as lead, zinc, and 
copper, tarnish or rust slowly ; i.e. they combine slowly with 
the oxygen of the air. The chemical activity of oxygen at 
high temperatures is readily shown by putting burning or 



OXYGEN — OZONE 2 1 

glowing substances into it. All burn vividly in oxygen. It 
interacts chemically with the substances, and the intensity 
of the chemical energy is shown by the heat and light evolved. 
The vigorous chemical activity of oxygen can be shown by 
simple experiments. When a glowing stick of w r ood is put 
into oxygen, the wood instantly bursts into flame ; and if 
left in the oxygen, the wood continues to burn brightly until 
the gas is exhausted. If faintly glowing charcoal is put into 
oxygen, the charcoal glows vividly. Sulphur burns in air 
with a feeble, blue flame, but in oxygen the flame is much 
larger and brighter. The flame in both cases is accompanied 
by a gaseous product which smells like a burning sulphur 
match. Iron wire does not burn in air, but if the end is 
coated with burning sulphur and then put into a bottle of 
oxygen, the wire burns vividly, throwing off a shower of 
sparks ; a globule of white-hot molten iron oxide is often seen 
on the end of the wire, and sometimes the inside of the 
bottle is coated with a reddish powder, which is mainly a 
compound of iron and oxygen. Iron and oxygen combine 
at a higher temperature than do sulphur and oxygen, so the 
heat from the burning sulphur is needed to start the chemi- 
cal action between the oxygen and the iron. When lighted 
magnesium is put into oxygen, the burning metal instantly 
becomes surrounded with a dazzling flame, and burns 
rapidly to a white powder, thus showing that the tempera- 
ture at which it begins to combine with oxygen is much lower 
than that in the case of iron. 

Test for Oxygen. — The conspicuous behavior of a glow- 
ing stick or burning substance when put into oxygen enables 
us to distinguish oxygen from all other gases except one 
(see Nitrous Oxide). Thus, if a glowing stick is thrust suc- 
cessively in three bottles of gas and relights in one, that gas 
is probably oxygen. This critical examination to establish 



22 INORGANIC CHEMISTRY 

the identity of oxygen is called testing or making a test. 
Each element has properties which respond to appropriate 
tests. All compounds likewise behave in some decisive way 
when subjected to tests. These tests have been established 
by extensive investigation and are extremely useful in chem- 
istry. 

Chemical Changes illustrated by the Chemical Properties 
of Oxygen. — In all the chemical changes described in the 
preceding paragraphs two features are conspicuous, viz. the 
loss of physical properties of the interacting elements and 
the formation of new substances. The carbon (of the wood 
and charcoal), iron, sulphur, and magnesium are no longer 
recognizable as such in the compounds produced by the 
chemical change. The oxygen likewise disappears; for 
although two products are gases, their properties are not 
those of oxygen. Thus, one gas (carbon dioxide) extinguishes 
flames and the other (sulphur dioxide) has a suffocating 
odor. The chemical change in the case of the four elements 
just mentioned is combination. The oxygen was added chem- 
ically to each element and the product was a compound of 
the two elements called an oxide. Their names are carbon 
dioxide, sulphur dioxide, ferrous-ferric oxide, and magnesium 
oxide. Any compound whose constituents are oxygen and 
one other element is an oxide. (See Oxides, below.) The 
chemical change in the case of wood is not simple combina- 
tion. Wood contains chiefly a compound of carbon, hydro- 
gen, and oxygen, which is first decomposed by heat, and then 
the carbon and the hydrogen combine individually with 
oxygen as in the other cases. 

Oxidation. — When sulphur, iron, magnesium, and carbon 
(in wood and charcoal), and other elements burn in oxygen, 
they combine rapidly with it. This change is oxidation, 



OXYGEN — OZONE 23 

Oxidation almost invariably produces heat, though in slow 
oxidation the heat is not readily detected and is sometimes 
overlooked. Very rapid oxidation is accompanied by light. 

The fact that burning in its simplest aspect is a combining 
with oxygen can be easily verified. It has been repeatedly 
shown that oxygen is one constituent of all the products 
formed by burning substances in that gas. Thus, carbon 
forms an invisible gas called carbon dioxide, w T hich is a com- 
pound of carbon and oxygen. Similarly, sulphur, iron, and 
magnesium form compounds of these elements and oxygen. 
These facts may be further verified by a simple experiment. 
If mercury is heated, it gains in weight, and red particles 
collect on its surface; but if it is protected from the air by 
some coating and then heated, there is no gain in weight 
and no evidence of the red product. Therefore, when the 
exposed mercury is heated, something from the air must be 
added to it. Now, if the red substance is collected and 
heated in a glass tube, mercury and oxygen are the only 
products. Hence, the exposed mercury, when heated, must 
have combined with the oxygen of the air. 

The process of oxidation is not always so simple as that 
described in the preceding paragraph. Often one or more 
of the constituents of a compound is seized upon by the 
oxygen, so to speak, and converted into an oxide. Hydro- 
gen is especially liable to be thus removed from such com- 
pounds as contain it. Hence oxidation, which is primarily 
a combining with oxygen, is often a complex process resulting 
in the decomposition of a compound and the formation of 
one or more oxides. 

Oxidation is not always rapid enough to produce light 
and appreciable heat. Thus, iron and other metals rust 
and wood decays slowly, but both processes are essentially 
oxidation. Sometimes slow T oxidation develops considerable 
heat. Thus, oily rags, piles of hay, and heaps of coal often 



24 INORGANIC CHEMISTRY 

take fire unexpectedly, because the heat produced by the 
continued oxidation cannot escape. Fires caused by such 
oxidation are sometimes said to be due to spontaneous 
combustion. 

Many substances give up oxygen readily and are called 
oxidizing agents. In this class belong potassium chlorate, 
which is used in fireworks, and potassium nitrate, which is 
an ingredient of gunpowder. In the process of oxidation, 
oxidizing agents lose oxygen, and are said to undergo reduc- 
tion ; the latter process will be more fully described in the 
next chapter. 

Oxides have already been defined as compounds of oxygen 
and one other element. There are many oxides, and their 
names often express in a general way their composition. 
Oxides of different elements are distinguished by placing the 
name of the element (or a slight modification of it) before the 
word oxide, e.g. magnesium oxide, lead oxide, zinc oxide, 
nitric oxide, ferric oxide. Sometimes di-, or a similar 
numerical syllable, is prefixed to the word oxide, e.g. carbon 
monoxide, manganese dioxide, barium dioxide, sulphur 
trioxide, phosphorus pentoxide. The prefix indicates the 
proportion of oxygen in the oxide. 

Combustion. — In ordinary language combustion means 
fire or burning. Substances which kindle and burn readily 
are called combustible, while those which burn with diffi- 
culty or not at all are called incombustible. In a limited 
sense, combustion is rapid oxidation accompanied by heat 
and often also by light. Oxygen is essential to ordinary 
combustion. If air is excluded from a fire, the fire goes out. 
When wood, coal, paper, oil, or any other combustible sub- 
stance burns, the carbon (of which they partly consist) unites 
with the oxygen of the air, thereby forming the invisible gas 
carbon dioxide, while the chemical change is attended by 



OXYGEN — OZONE 25 

heat and light. Briefly, a burning substance is uniting rap- 
idly with oxygen. But since air is only about one fifth oxy- 
gen (the remainder being chiefly nitrogen, which does not 
support combustion), combustion is less rapid and hence 
less vigorous in air than in oxygen. The compounds formed 
during combustion are oxides; in the case of fuel they are 
usually carbon dioxide and water — the latter being hydro- 
gen oxide. The temperature at which combustion takes 
place varies between wide limits. Some substances, like 
phosphorus and gasolene vapor, catch fire at a moderate 
temperature, while others do not burn until heated to the 
very highest temperatures. Each substance has its own 
kindling temperature, i.e. the temperature to which it must 
be heated before it will catch fire, though this temperature 
depends on the form of the substance as well as on its nature 
(see page 172). Application of this fact is seen in the use 
of paper and kindling wood in starting a fire in a stove. 

The correct explanation of fire, burning, and combustion was first 
made by Lavoisier (1743-1794). For many years chemists had be- 
lieved that all combustible substances contained a principle called 
phlogiston, and that when a substance burned, phlogiston escaped. 
Very combustible substances were thought to contain much phlogis- 
ton, and incombustible substances no phlogiston. This theory of 
combustion was proposed by Becher (1635-1682) and advanced by 
Stahl (1660-1734). Many famous chemists — Priestley, Scheele, and 
Cavendish — supported it. Lavoisier, in 1775, proved by his own 
and others' experiments that phlogiston did not exist, and that 
ordinary combustion is a process of combination with "a certain 
substance contained in the air." Soon after he identified this sub- 
stance as oxygen. 

Combustion, in a broad sense, is not exclusively oxidation, but any 
chemical change which is attended by light and heat. Thus, hydro- 
gen and other elements burn in chlorine, and phosphorus often pro- 
duces heat and light by its combinations with other elements than 
oxygen. This broader meaning will be illustrated and discussed 
later. 



26 INORGANIC CHEMISTRY 

Relation of Oxygen to Life. — Oxygen is essential to all 
forms of animal and plant life. If an animal or a plant is 
deprived of air, it dies. By respiration air is drawn into the 
lungs, where part of its oxygen is given to the blood ; this oxy- 
gen, which is distributed to all parts of the body by the blood, 
oxidizes the tissues of the body. By this slow oxidation 
waste products are formed and heat is supplied to the body. 
One of these waste products is carbon dioxide gas, which 
with other gases is exhaled from the lungs. New tissue is 
built up from the food we eat. The human body resembles 
a steam engine. In each, the oxygen of the air helps burn 
fuel largely composed of carbon. In the engine, the products 
escape through a chimney and the heat produced is used to 
form steam which moves parts of the machine; in the body, 
the products escape mainly through the lungs and the heat 
keeps the body at the temperature at which it can best 
perform its functions. 

The vital necessity to fish of the small proportion of oxygen 
dissolved by water should be recalled in this connection. 

Oxygen is often administered to a person who is too weak 
to inhale enough air. Oxygen is also used in various forms 
of emergency breathing apparatus. The pulmotor, for 
example, is a kind of pump by which air rich in oxygen can 
be forced into the lungs at the normal rate of respiration ; 
it is used to resuscitate persons who have been overcome by 
smoke or gas or who have been rendered unconscious by an 
electric shock. An oxygen-breathing apparatus is used by 
men who must work in places containing smoke or poisonous 
gases, e.g. in a mine after an explosion. The apparatus con- 
sists essentially of a helmet or a mouth-breathing device con- 
nected by flexible tubes with a cylinder of compressed oxygen 
and a regenerating can. The apparatus is carried much like 
a knapsack. The supply of oxygen is regulated by a valve 
in the cylinder, the nitrogen in the original supply of air is 



OXYGEN — OZONE 27 

breathed over and over, and the carbon dioxide and water 
vapor from the lungs are absorbed by potassium hydroxide 
in the regenerating can. A trained man equipped with this 
apparatus can work as long as two hours in a vitiated atmos- 
phere. 

Decay of organic matter involves oxidation, which is often 
hastened by bacteria. The carbon and hydrogen in the 
organic matter form principally carbon dioxide and water; 
decay in this respect resembles combustion and respiration. 

Uses of Oxygen. — Oxygen for commercial use is stored 
under pressure in strong iron cylinders. A mixture of oxy- 
gen and hydrogen gas or acetylene gas if burned in a suitable 
apparatus produces an intensely hot flame. The oxy- 
hydrogen flame is used to melt refractory metals and to pro- 
duce the brilliant light of the stereopticon ; the acetylene 
flame is used in welding and also in dismantling iron and steel 
structures, e.g. bridges, fences, and frames of buildings. 

Discovery of Oxygen. — Oxygen was discovered on August 
1, 1774, by Priestley (1733-1804). He prepared it by focus- 
ing the sun's rays upon red mercuric oxide by means of "a 
burning lens of twelve inches focal distance. " It was inde- 
pendently discovered by Scheele (1742-1786), a Swedish 
chemist, about the same time. 

Priestley called the gas dephlogisticated air because he 
regarded it as " devoid of phlogiston." Scheele called it 
empyreal air, i.e. fire air or fire-supporting air, because it 
assisted combustion. Lavoisier, in 1778, gave it the name 
oxygen (from the Greek oxus, acid, and gen, the root of a verb 
meaning to produce) because he believed from his experiments 
that oxygen was necessary for the production of acids — a 
view now known to be incorrect. 



28. INORGANIC CHEMISTRY 

Problems 

1. Potassium chlorate contains 39.18 per cent of oxygen. How 
many grams of oxygen can be prepared from 725 gm. of potassium 
chlorate ? 

2. What approximate weight of oxygen can be prepared from 100 
gm. of potassium chlorate containing 12 per cent of impurity? 

3. What is the weight of (a) 10 1. of oxygen, (b) 75 1., (c) 500 cc, 
(d) 750 cc, (e) 4 1. (at standard temperature and pressure) ? 

4. A room 25 m. long, 17 m. wide, and 15 m. high is filled with 
oxygen. What weight of gas does it contain? 

5. Perform the problems in the Appendix, § 2. 



Ozone 

Ozone is a gas related to oxygen, though its properties 
differ. 

Formation and Preparation. — Ozone is formed when 
electric sparks pass through the air, and is therefore pro- 
duced when electrical machines are in operation and during 
thunder storms. Slow oxidation, especially of moist phos- 
phorus, produces ozone. Its formation accompanies the 
burning of hydrogen (in oxygen), and the passage of elec- 
tricity through a solution of sulphuric acid and water. 
Ozone is prepared by subjecting cold, dry oxygen or air to a 
silent electric discharge. 

Properties. — Ozone has a peculiar odor, suggesting burn- 
ing sulphur. The name ozone signifies smell. It has a 
bluish color, and at a low temperature condenses to a blue 
liquid. Liquid ozone boils at — 119° C. under atmospheric 
pressure. It is a powerful oxidizing agent, tarnishing metals, 
bleaching colored vegetable substances, deodorizing foul 
animal matter, and corroding such substances as cork and 
rubber. When heated to 250° C, or higher, it is wholly 
changed into oxygen. When oxygen is changed into ozone, 



OXYGEN — OZONE 29 

it is found that three volumes of oxygen yield two volumes 
of ozone ; and, conversely, the two volumes of ozone, when 
heated, become three volumes of oxygen. Hence, volume 
for volume, ozone is 1.5 times heavier than oxygen. Its 
theoretical relation to oxygen will be subsequently discussed. 

Uses. — Pure ozone is seldom prepared, but air containing 
ozone is used to sterilize drinking water, as a bleaching agent, 
and a disinfectant. 

Exercises and Problems 

1. Suggest experiments to show (a) that in preparing oxygen 
from a mixture of potassium chlorate and manganese dioxide the 
latter is unchanged, and (6) that air contains oxygen. 

2. Define and illustrate (a) oxide, (b) oxidation, (c) oxidizing 
agent. 

3. Cite cases of spontaneous combustion. 

4. Potassium chlorate contains 39.18 per cent of oxygen. How 
many grams are needed to prepare 25 1. of oxygen (at 0° C. and 
760 mm.)? 

5. Potassium chlorate contains 39.18 per cent of oxygen. 
How many liters of oxygen (at 0° C. and 760 mm.) can be prepared 
from 75 gm. of potassium chlorate ? 

6. Mercuric oxide contains 7.4 per cent of oxygen. What weight 
of mercuric oxide must be decomposed to yield 25 gm. of oxygen ? 

7. Mercuric oxide contains 7.4 per cent of oxygen. How 
many grams of mercuric oxide will be needed to prepare 5 1. of 
oxygen gas (at 0° C. and 760 mm.) ? 

8. Mercuric oxide contains 7.4 per cent of oxygen. How many 
liters of oxygen gas (at 0° C. and 760 mm.) can be prepared from 
525 gm. of mercuric oxide ? 

9. If one third of the oxygen in manganese dioxide can be liber- 
ated by heating, what (a) weight and (b) volume of oxygen can be 
obtained from 25 gm. of manganese dioxide ? (Note : Manganese 
dioxide contains 36.78 per cent of oxygen.) 

10. When 2 gm. of a certain substance were heated, all the 
oxygen which the substance contained was given off, and a residue 
weighing 1.07 gm. was left. Calculate the percentage of oxygen 
in the substance. 



CHAPTER III 
Hydrogen 

Occurrence. — Free hydrogen gas is present in the gases 
which escape from volcanoes, petroleum wells, and natural 
gas openings. Artificial illuminating gas contains consider- 
able hydrogen. It is a product of fermentation and decay, 
and according to recent observations a very small quantity 
is present in the atmosphere of the earth. Enormous quan- 
tities of free hydrogen exist in the atmosphere of the sun, and 
during an eclipse of the sun gigantic streams of glowing hy- 
drogen may be seen shooting out from the sun's disk thou- 
sands of miles into space. Other heavenly bodies contain 
hydrogen. Meteorites, which come from regions far beyond 
our earth, sometimes contain free hydrogen. 

Combined hydrogen is abundant and widely distributed. 
It forms about one ninth (by weight) of water. Animal and 
vegetable substances contain hydrogen in combination with 
oxygen and carbon, and sometimes nitrogen. It is an essen- 
tial constituent of all acids. Combined with carbon, it 
forms many gases and liquids called hydrocarbons, which 
are constituents of illuminating gas, kerosene, and gasolene. 
Combined with carbon and oxygen, it forms such compounds 
as sugar, starch, paper, wood, and numerous artificial prod- 
ucts. With nitrogen it forms the familiar compound, am- 
monia, and with sulphur, the bad smelling gas, hydrogen 
sulphide, which occurs in many sulphur springs. 

Preparation. — Hydrogen, like oxygen, is prepared from 
its compounds. This is usually done by allowing certain 

30 



HYDROGEN 31 

metals and acids to interact. The metals usually employed 
are zinc, iron, or magnesium, and the acids are dilute sul- 
phuric acid or hydrochloric acid. The hydrogen comes from 
the acid and bubbles through the liquid, when the acid and 
metal are mixed. Hydrogen is prepared in the laboratory 
by putting the metal and dilute acid in a glass vessel provided 
with a delivery tube arranged to collect gas over water in a 
pneumatic trough. No flame should be near during the 
performance of this experiment, because a mixture of air and 
hydrogen explodes violently when ignited. The interaction 
of zinc and sulphuric acid produces, besides hydrogen, a 
soluble compound called zinc sulphate. 

Hydrogen can also be prepared from bases (compounds of 
hydrogen, oxygen, and a metal). Thus, when aluminium 
is boiled with sodium hydroxide, hydrogen is formed. 

Hydrogen can be obtained from water by allowing sodium 
or potassium to react with it. If a small piece of sodium is 
dropped upon cold water, the sodium melts into a shining 
globule, which spins about rapidly on the surface with a 
hissing sound, and finally disappears after a slight explosion. 
But when the sodium is wrapped in a piece of tea lead pierced 
with a few holes and then dropped beneath the shelf of a 
pneumatic trough filled with water, the action proceeds 
smoothly; hydrogen gas rises and displaces the water from 
a test tube or bottle supported over the hole in the shelf. 

Hydrogen can also be prepared by the interaction of steam 
— the gaseous form of water — and certain other metals, 
if they are heated. This experiment was first performed by 
Lavoisier, in 1783, while he was studying the composition of 
water. He passed steam through a red-hot gun barrel con- 
taining bits of iron, and hydrogen escaped from the tube. 
Since Lavoisier was studying the composition of water and 
not the properties of hydrogen, he naturally thought of this 
gas as essential for forming water. So he says in his notes, 



32 INORGANIC CHEMISTRY 

"No name appears to us more suitable than that of hydro- 
gen, that is to say, 'generative principle of water.'" Apart 
from historical interest, this experiment has commercial im- 
portance, for if steam is passed over white-hot coal (instead 
of iron), producer gas is formed. This mixture consists of 
about one-half hydrogen, and is used as a source of heat in 
making steel and glass. If oil vapor is added to this mixture, 
water gas is formed. This is an illuminating gas, and is used 
in many cities. (See Water Gas.) 

Hydrogen, together with oxygen, is liberated from water 
by passing a current of electricity through water containing 
a little sulphuric acid. (See Chapter V.) 

Chemical Changes illustrated by the Preparation of Hy- 
drogen. — The preparation of hydrogen by the interaction 
of a metal and an acid illustrates the third kind of chemical 
change, viz. substitution, or, as it is sometimes called, dis- 
placement or replacement. Zinc and dilute sulphuric acid 
are usually used. The hydrogen is displaced from the acid 
and the zinc takes its place; i.e. zinc is substituted chemi- 
cally for hydrogen. The chemical change is not essentially 
different from decomposition and combination, for we might- 
picture the acid as decomposing into two fundamental parts, 
one part escaping as hydrogen gas, the other part (the sul- 
phur and oxygen) at once combining with the zinc to form 
the new compound, zinc sulphate. It is better, however, to 
regard substitution as a chemical change in which one ele- 
ment replaces another in a compound, thereby producing a 
different element and a different compound. The chemical 
change in the preparation of hydrogen from zinc and sul- 
phuric acid can be expressed by the following equation: — 

Zinc + Sulphuric Acid = Hydrogen + Zinc Sulphate. 

The preparation of hydrogen by the interaction of water and 
sodium is also a case of substitution. Here the sodium dis- 



HYDROGEN 33 

places one half of the hydrogen of the water, thereby pro- 
ducing free hydrogen and a new compound (sodium hydrox- 
ide), which consists of sodium, oxygen, and the rest of the 
hydrogen. The sodium hydroxide dissolves in the water in 
the trough. Briefly, sodium is substituted for hydrogen. 
The change can be expressed thus : — 

Sodium + Water = Hydrogen + Sodium Hydroxide. 

Physical Properties. — Hydrogen has no taste or color. 
The pure gas has no odor, though hydrogen as ordinarily 
prepared may have a disagreeable odor, due to impurities. 
Hydrogen is the lightest known substance. 
One liter of dry hydrogen at 0° C. and 760 mm. 
weighs only 0.08987 gm. Volume for volume, 
hydrogen is about 14.4 times lighter than air 
and about 16 times lighter than oxygen. 

The extreme lightness of hydrogen can be 
easily shown. (1) If a wide-mouth bottle of 
the gas is left uncovered two or three minutes 
and a lighted match then dropped in, the match 
will continue to burn. If hydrogen had been 
present, it would have combined with the Fig. 2. — Pour- 
oxygen of the air with a loud explosion, as mg hydrogen 

in r ^ iiii upward. 

soon as the name of the match reached the 
mixture. (2) If a bottle of hydrogen is held beneath a bottle 
of air in the position shown in Figure 2, the gases exchange 
places, the hydrogen, owing to its lightness, rising into the 
upper bottle. Its presence there can be readily shown by 
inserting a lighted match into this bottle; if the experiment 
has been properly done, the hydrogen will burn quickly and 
quietly, but in most cases a loud explosion shows that only 
a part of the hydrogen has succeeded in entering the upper 
bottle. A lighted match dropped into the other bottle re- 
veals air, or a mixture of air and hydrogen (if the experiment 




34 INORGANIC CHEMISTRY 

has been performed too hastily). (3) If a small collodion, 
or rubber, balloon is filled with hydrogen and then released, 
it will rise rapidly into the air. Hydrogen, because of its 
lightness, is used to fill balloons; but ordinary illuminating 
gas, which is cheaper, is sometimes used. 

Hydrogen is not very soluble in water, but it is absorbed 
by several metals, especially the rare metal palladium. The 
absorption of gases by metals is called occlusion. Only 
about 1.84 1. of hydrogen at 760 mm. pressure dissolve in 
100 1. of water at 20° C. Palladium absorbs from 300 to 900 
times its own volume of hydrogen, according to the condi- 
tions of the experiment. Platinum, gold, and iron act simi- 
larly, though to a less degree. Illuminating gas, which con- 
tains considerable hydrogen, is also absorbed by certain 
metals. Heat is developed by occlusion. This heat may be 
sufficient to raise the metal to a red heat and to ignite the gas 
itself. One form of self-lighting gas burner acts on this 
principle. Occlusion is partly chemical and partly physical. 

Hydrogen diffuses readily; i.e. it quickly passes through 
porous substances, mixes with other gases without stirring 
or agitating, and freely escapes into space in all directions. 
Hydrogen has the highest rate of diffusion, because its den- 
sity is the lowest, it being a general fact that the rate of dif- 
fusion of a gas is inversely proportional to the square root of 
the density. Thus, the rate of diffusion of hydrogen is 
four times that of oxygen, since the density of oxygen is six- 
teen times that of hydrogen. We are largely indebted for 
our knowledge of diffusion to the English chemist, Thomas 
Graham (1805-1869). 

Liquid hydrogen is colorless and transparent. It was first 
obtained by Dewar in 1898. At the ordinary pressure of the 
atmosphere liquid hydrogen boils at — 252.5° C. When 
cooled to about — 256° C. by evaporation under reduced pres- 
sure, the liquid becomes a mass of solid hydrogen; the latter 




SIR JAMES DEWAR 



HYDROGEN 35 

is a white froth if produced while boiling, and a transparent, 
ice-like solid if cooled w T hen quiet. Solid hydrogen melts at 
— 260° C. (if the pressure is 58 mm.). 

Hydrogen is not poisonous, if pure. It does not support 
life, but a little may be breathed without danger. When 
the lungs are filled with it, the voice becomes very thin and 
shrill. 

Chemical Properties. — Hydrogen combines readily with sev- 
eral elements, especially chlorine, oxygen, sulphur, and nitro- 
gen. The conditions favorable for combination vary greatly. 

When organic matter containing nitrogen and sulphur 
decays, the products include ammonia and hydrogen sul- 
phide. The former is a compound of nitrogen and hydrogen, 
and the latter of sulphur and hydrogen. Hydrogen and 
chlorine gases do not unite in the dark, but in the sunlight 
they combine with explosive violence even at the ordinary 
temperature. Hydrogen gas burns quietly in chlorine 
gas. The flame is bluish white, not very hot, and the 
product is hydrochloric acid gas — a compound of hydro- 
gen and chlorine. This burning of hydrogen and chlorine 
illustrates the broader use of the word combustion. No 
oxygen is involved. It is a chemical change attended by 
light and heat and belongs to the class of changes called 
combination, since the hydrogen and chlorine are added 
or combined chemically. 

Hydrogen and oxygen do not unite at the ordinary tempera- 
ture, but at about 750° C. the gases combine with explosive 
violence. This temperature is provided by an ordinary 
flame or red-hot wire. A mixture of hydrogen and air 
explodes when ignited. Therefore, the air should be fully 
expelled from the apparatus in which hydrogen is being 
generated before the gas is collected, and no flame, large 
or small, should be near. Neglect of these precautions 



36 



INORGANIC CHEMISTRY 



has caused serious accidents. A small jet of hydrogen, 
however, burns quietly in air or oxygen. The flame is 
almost invisible and very hot. Water is the sole product 
of the combustion of hydrogen. The equation for this 
chemical change is : — 

Hydrogen + Oxygen = Water. . 

These properties of the hydrogen flame can be readily 
shown by generating hydrogen in a suitable apparatus 
and lighting the dry gas as it issues from a small opening 
(Fig. 3). A platinum wire quickly becomes red hot in the 



« 




Fig. 3. — Apparatus for burning hydrogen. Acid is slowly introduced 
through the funnel into the flask, which contains zinc. The liberated hy- 
drogen is dried as it passes through the U-tube containing calcium chloride 
and is lighted at the platinum tip after all the air has been driven from 
the apparatus. The tip, which is attached to the delivery tube by a small 
rubber tube, is shown (about actual size) on the left. 



flame, and a dry inverted bottle into which the flame 
is inserted becomes coated on the inside with moisture 
condensed from the steam. The film of water often noticed 
on the bottom of a vessel placed over a lighted gas range 
or a Bunsen burner is formed by the burning of the hydrogen 
and of the hydrogen compounds in the illuminating gas. 
The fact that water is the only product of burning hydro- 
gen was first shown in 1783 by Cavendish (1730-1810). 



HYDROGEN 37 

Lavoisier in the same year verified this fact and utilized 
it to explain the composition of water. 

The temperature of the hydrogen flame is very high.. 
More heat is produced by burning hydrogen and oxygen 
than by burning the same weight of any other substance. 

Hydrogen does not support combustion, as the term is 
usually used. This property is illustrated by putting 
a lighted taper into an inverted bottle of hydrogen. The 
taper ignites the hydrogen, which burns at the mouth of the 
bottle, but the taper does not burn inside the bottle. Hence, 
hydrogen burns, but does not support combustion. When 
the extinguished taper is slowly withdrawn through the 
burning hydrogen, however, it is relighted. 

Hydrogen removes oxygen from compounds (see Reduction). 

Test for Hydrogen. — The test for hydrogen is that it 
extinguishes a flaming stick but is lighted at the same time, 
often with an explosion, and continues to burn until the 
gas is exhausted. 

The Oxyhydrogen Blowpipe utilizes the intense heat 
produced by burning a mixture of hydrogen and oxygen. 
The apparatus (Fig. 4) 
consists of two pointed 
metal tubes. The inner 
and smaller one is for the 

oxygen, and the Outer ^ 4. ~ Oxyhydrogen blowpipe tip. 

and larger one for the hydrogen. Their pointed ends are 
close together, and the two gases mix as they are forced 
out of these small openings by the pressure maintained in 
the storage tanks. Sometimes the tubes are separated, 
but the gases flow from a similar opening. The hydrogen 
is first turned on and lighted at the pointed opening, then 
the oxygen is turned on and the flow gradually regulated 
until the flame is the desired size, usually thin, straight, 




38 



INORGANIC CHEMISTRY 



and as long as required. There is no danger in using the 
blowpipe, provided it does not leak and the pressure of the 
gases is properly regulated by the stopcocks. In the hot 
flame, some metals, like silver, turn to vapor; some, like 
iron, burn brilliantly; while others, like platinum, melt. 
When the flame strikes against a piece of lime, the latter 
becomes intensely bright. Thus used, it is called the lime 
or calcium light, and is utilized in the stereopticon. The tem- 
perature of the oxy hydrogen flame is from 1800 to 2500° C. 

The blast lamp is a modification of the oxyhydrogen 
blowpipe. The apparatus (Fig. 5) consists of two tubes, 

an inner one for air and an outer 
one for illuminating gas. The air, 
which is forced through the apparatus 
by a bellows, provides oxygen, w T hile 
the illuminating gas contains hydro- 
gen and other combustible gases. 
The mixture burns at the opening 
of the tubes with a colorless or bluish 
flame, which is hotter than the Bun- 
sen flame (the usual source of heat 
for chemical experiments) . The shape 
and size of the flame are easily regu- 
lated by stopcocks. The blast lamp is used as a source of 
heat for many operations in the laboratory, especially in 
chemical analysis. 




Fig. 5. — Blast lamp. 



Reduction. — Hydrogen not only combines energetically 
with free oxygen, but it also withdraws oxygen from com- 
pounds. The chemical removal of oxygen is called re- 
duction. Hydrogen is a vigorous reducing agent, just as 
oxygen is an energetic oxidizing agent. When oxides of 
certain metals are heated in a current of hydrogen, the oxy- 
gen of the oxide is removed and combines with the hydrogen 



HYDROGEN 39 

to form water; the metal is left uncombined. Thus, by 
heating lead oxide in hydrogen, water and metallic lead are 
produced. Chemically speaking, the lead oxide is reduced 
by the hydrogen. The chemical change is substitution 
(the hydrogen being substituted chemically for the metal), 
and it can be expressed thus : — 

Hydrogen + Lead Oxide = Water + Lead. 

This chemical change can also be interpreted from the 
standpoint of oxidation, because the hydrogen is oxidized 
to water at the same time the lead oxide is reduced. In 
fact, the processes of reduction and oxidation are closely 
related and usually occur in the same chemical change; 
either one may be emphasized in interpreting the change. 
It is preferable, however, at this stage to define reduction 
as the removal of oxygen from a compound, postponing the 
details of the process until more facts are available. In its 
simplest form, reduction is the opposite of oxidation. ' 

Discovery of Hydrogen. — Paracelsus in the sixteenth 
century obtained hydrogen by the interaction of acids and 
metals. It was identified as an element in 1766 by Caven- 
dish, who called it inflammable air. The name hydrogen, 
given to it by Lavoisier, in 1783, is derived from the Greek 
word hudor, water, and gen, the root of a verb meaning to 
produce. (See Preparation of Hydrogen, third paragraph.) 

Problems 

1. What volume does 5 gm. of hydrogen occupy (at 0° C. and 
760 mm.) ? 

2. The density of chlorine is nearly thirty-six times that of hydro- 
gen. Compare its rate of diffusion with that of hydrogen. 

3. A certain gas passes through a porous partition 2.5 times 
slower than hydrogen. What is its density? 

4. How many times heavier than a liter of hydrogen is a liter of 
oxygen, both being dry and under standard conditions? 



40 INORGANIC CHEMISTRY 

5. What is the weight of (a) 8000 cc. of dry hydrogen gas at 0° 
C. and 760 mm. ? (6) Of 1800 cc. ? (c) Of 9 1. ? 

6. The standard pressure at which a gas is measured is 760 
mm. Express the same in inches. 

7. If sulphuric acid contains 2.04 per cent of hydrogen, how 
many liters of hydrogen (at 0° C. and 760 mm.) can be obtained 
from 137 gm. of sulphuric acid? 

8. Water contains 11.18 per cent of hydrogen. How many 
(a) cc. and (b) gm. of the gas can be prepared from 1 1. of water? 

9. What is the weight in gm. of (a) 50 1. of hydrogen gas? 
(6) 50,000 cc. ? (c) 50 cdm. ? 

10. (a) 72 1. of hydrogen gas (at 0° C. and 760 mm.) weigh 
how many grams, (b) centigrams, (c) milligrams, (d) kilograms, 
(e) decigrams? 

11. How many grams does a cubic meter of hydrogen weigh? 

12. (a) How many cc. (at 0° C. and 760 mm.) will 75 gm. of 
hydrogen occupy? (b) How many gm. will 75 cc. weigh? 

13. A cylindrical tank 1.5 m. long and 30 cm. in diameter is 
filled with hydrogen (at 0° C. and 760 mm.), (a) What is the 
weight of the gas in gm. ? (b) How many gm. of water is needed 
to prepare the gas? (c) How many kg. of sulphuric acid? (d) 
How many dg. of sodium hydroxide (containing 1.75 per cent of 
hydrogen) ? 



CHAPTER IV 
Some Properties of Gases 

Introduction. — Several elements and many compounds 
are gases or can be readily changed into gases. Elementary 
gases, besides oxygen and hydrogen, are chlorine and nitrogen; 
compound gases are carbon dioxide, hydrochloric acid, 
ammonia, and sulphur dioxide. Air is a mixture of several 
gases, but it behaves like a single gas. Water is readily con- 
verted into the gaseous state, which is familiar as steam and 
water vapor. 

The properties of all gases vary with the temperature 
and the pressure to which they are subjected. Thus, 
oxygen gas becomes liquid oxygen at a low temperature, 
while liquid water is constantly changing into water vapor. 
The most common and conspicuous change, however, is an 
alteration in volume whenever pressure and temperature 
are varied. 

Relation of Gas Volumes to Temperature and Pressure. — 

The actual volume occupied by a gas depends upon the 
temperature and pressure prevailing at the time of obser- 
vation. The volume expands with rise of temperature or 
with decrease of pressure; it contracts with fall of tempera- 
ture or with increase of pressure. In general, if we cool a 
gas or subject it to a greater pressure, it shrinks ; if we heat a 
gas or subject it to a lower pressure, it expands. By common 
consent, the normal or standard temperature is zero degrees 
on the centigrade thermometer (or briefly 0° C), and the 
normal or standard pressure is the pressure indicated by the 

41 



42 INORGANIC CHEMISTRY 

barometer when the mercury column is 760 millimeters high 
(or briefly 760 mm.). Under these conditions, which are 
called standard conditions, a liter of dry oxygen gas weighs 
1.429 gm. But at another temperature or pressure the 
liter would contain a different quantity of oxygen gas, and 
would therefore have a different weight. For example, if 
the pressure is increased, the volume becomes less, more gas 
must be added to bring the volume up to a liter, and this 
second liter of oxygen would weigh more than 1.429 gm. 
That is, a liter vessel, when full, always contains a liter, 
but the weight of the contents varies with the quantity of 
gas contained in this volume. Clearly, if we wish to com- 
pare the weights of gases by means of their volumes, we must 
know the conditions under which the volume is measured. 
As we shall subsequently see, the comparison of weights of 
gases is a frequent and highly important operation in chemis- 
try. Some method is necessary, therefore, to permit this 
comparison. If all gases could be measured at 0° C. and 
760 mm., their volumes would be comparable and the 
weights deduced or obtained directly from these volumes 
would be a true measure of the actual quantity of the gases 
in the observed volumes. But it is experimentally incon- 
venient to measure gases at 0° C. and 760 mm. So. it is 
customary to measure the volume under the conditions 
existing at the time of the experiment, and then reduce the 
observed volume to the volume it would occupy under stand- 
ard conditions. This mathematical reduction is performed 
by applying two laws, — the law of Charles and the law of 
Boyle. 

Relation of the Volume of a Gas to Changes in Tempera- 
ture. — It has been found by experiment that all gases 
under constant pressure expand or contract equally for 
equal changes of temperature. If the volume of a gas is 



SOME PROPERTIES OF GASES 



43 



measured at 0° C, the gas expands or contracts 1/273 of this 
volume for a rise or fall of one degree. That is, 273 vol- 
umes at 0° C, become 274 at 1° C., 275 at 2°, 280 at 7°, 272 
at -1°, 270 at -3°, or in general, (273 + volumes at 
f {i.e. at any temperature). The statement above summa- 
rizing the relation between gas volumes and temperature 
is known as the Law of Charles. It applies accurately 



125 






400 






100 




- 


375 
373 






75 




— 


350 






50 




— 


325 


Centigrade 


Absolute 


25 






300 


200 


473 







- 


275 

273 


100 


373 


3-25 






250 „ 


50 


323 


8 

02-50 




- 


225 | 


25 


298 


■d 






-^ 






£-75 
to 






200 # 

o 





273 


£-100 
5 






175 g 


-25 


248 


°-125 






150 


-50 


223 


-150 






125 


-100 


173 


-175 


— 




100 






-200 






75 


-273 





-225 






50 






-250 






25 






-273 














Fig. 6. — Centigrade and absolute thermometer scales (left) and some 
equivalent degrees (right). 



only to the temperature which would ordinarily be used, not 
to extreme temperatures nor to temperatures near the point 
at which the gas liquefies. Thus, according to the law, if 
a gas could be cooled to — 273° C, its volume would become 
zero ! But this low temperature has never been reached, 
and even if it could be, all gases (except possibly helium) 



44 INORGANIC CHEMISTRY 

become solid before reaching this temperature. Similarly, 
many gases dissociate or decompose at very high tempera- 
tures, and all deviate from the law when about to change 
from the gaseous to the liquid state. This point (— 273° C.) 
on the thermometer is called absolute zero, and a scale 
starting at 273° below 0° C. is called the absolute thermom- 
eter scale. The relation between the two scales is shown 
in Figure 6. 

An examination of these scales shows that absolute 
degrees are numerically greater by 273 than the correspond- 
ing centigrade degrees. This relation is sometimes expressed 
by the statement : — 

To convert centigrade into absolute add 273. 

The application of the law of Charles to the reduction 
of a gas volume to the volume it would occupy at 0° C. can 
be readily understood by an example. Suppose 10 cc. 
of dry oxygen gas at 15° C. are to be reduced to the volume 
occupied at 0° C. The corrected volume, as it is often called, 
can be found by two processes. (1) The first method utilizes 
the ordinary centigrade scale. Let the volume at 0° C. be 
represented by 273. Now since the volume of a gas at 0° C. 
expands 1/273 for each degree through which it is heated, 
the volume of oxygen at 15° C. would be represented by 
273 + 15. But 273 + 15 and 273 are in the same ratio as 10 
(the known volume at 15° C.) and X (the unknown volume 
at 0° C). Therefore we can state these relations in a 
proportion, thus : — 

273 + 15: 273:: 10 :X; X = 9.479 cc. 

Therefore, 10 cc. of oxygen at 15° C. occupy 9.479 cc. at 
0° C. Since t can be substituted for any temperature 
(above or below 0° C), the general form of the proportion 
can be written : — 

273 + 1 : 273 : : known vol. : vol. at 0° C. 



SOME PROPERTIES OF GASES 



45 



(2) The second method uses the absolute scale and is 
based on the fact that the volumes of a gas at different 
temperatures vary as their absolute temperatures. Suppose 
we have 273 cc. of a gas at 0° C. Since it expands 1/273 
of its volume at 0° C. for each degree of increase in tem- 
perature, and contracts 1/273 for each degree of decrease, 
its volumes would be as follows for certain centigrade 
temperatures : — 



Volume in cc. 


Centigrade Temperatures 


373 


100 


323 


50 


273 





223 


-50 


173 


-100 



Comparing these values with the corresponding absolute 
temperatures, we have the following relations : — 



Volume in cc. 


Centigrade Temperature 


Absolute Temperature 


373 


100 


373 


323 


50 


323 


273 





273 


223 


-50 


223 


173 


-100 


173 



It is clear from the first and third columns that the volumes 
and absolute temperatures are numerically the same. Hence 
in reducing gas volumes to those occupied at 0° C. by the 
absolute method, there are three steps, (a) Convert the 
observed centigrade temperature into absolute temperature 
by adding 273. (b) Make 273 the numerator of a fraction 
and the sum found in (a) its denominator. (c) Multiply 



46 INORGANIC CHEMISTRY 

the given or observed volume by the fraction formed in (b). 
Using the problem given above, the absolute process becomes 
(a) 273 + 15 = 288, (6) 273/288, (c) 10x273/288 = 9.479. 
The solution of a few problems by either of these methods 
will fix in mind the important relation of gas volumes to 
changes in temperature. 

Relation of Gas Volumes to Changes in Pressure. — It 
has been found, by experiment, that the volume of a gas at a 
constant temperature is inversely proportional to the pres- 
sure. This statement is the Law of Boyle, and was an- 
nounced by him in 1660. This law, like the law of Charles, 
applies accurately only to the pressures ordinarily used, but 
the general numerical relation between volume and pressure 
as stated above is sufficiently accurate for most purposes. 
Boyle's law means that the greater the pressure, the less the 
volume, and vice versa. The normal or standard pressure 
is the mean atmospheric pressure; this is equal to the 
pressure of a column of mercury 760 millimeters high. 
Briefly, the normal pressure is 760 mm. Most gases under 
examination are confined over water or some other liquid 
whose surface is exposed to the atmosphere, and since 
atmospheric pressure is transmitted through the liquid to 
the gas, the pressure which the gas is under is found by 
reading the pressure recorded by the barometer at the time 
the gas volume is read. The reduction of the observed 
volume to the volume it would occupy at 760 mm. is per- 
formed by applying Boyle's law. An actual case will make 
the process clear. Suppose we have 25 cc. of dry oxygen 
gas at 775 mm. and wish to know its volume at 760 mm. Two 
processes — not essentially different — can be used. (1) Ac- 
cording to Boyle's law, gas volumes are inversely pro- 
portional to the pressures; i.e. the observed pressure bears 
the same relation to the normal pressure as the normal 



SOME PROPERTIES OF GASES 47 

volume bears to the observed volume. This general relation 
applied to the problem becomes — 

775:760::X:25; X = 25.49. 

Therefore, 25 cc. of oxygen at 775 mm. occupy 25.49 cc. at 
760 mm. (2) The other method involves two steps, (a) Make 
760 the denominator of a fraction and the observed pressure 
its numerator, (b) Multiply the observed volume by the 
fraction formed in (a). Using the same problem, the second 
process becomes — 

(a) 775/760, (b) 25x775/760 = 25.49. 

Behavior of Gas Volumes under Simultaneous Action of 
Heat and Pressure. — Heat and pressure act independently 
upon a gas. That is, it is immaterial whether a gas is sub- 
jected to heat and to pressure at the same time or in succes- 
sion. The final volumes are equal in either case. Since both 
heat and pressure are factors in the cause of the changing 
gas volume, it is convenient to reduce a gas volume to the 
standard volume, so to speak, by a single operation. Thus, 
if the observed volume is 25 cc, the temperature 15° C, 
and the pressure 775 mm., the standard volume is found 
thus : — 

25 X 775/760 X 273/288 = X. 

Sometimes the reduction to standard conditions is per- 
formed by substituting the observed values in the formula: 

v== Y1P1 

760 [1 + (.00366 x t)] 

In the formula which merely involves the mathematical 
operations just described, V = final volume, V = observed 
volume, P' = observed pressure, t = observed temperature, 
and .00366 = 1/273. 1 

1 The method of deducing this formula is given in the author's " Experi- 
mental Chemistry," pp. 361-363. 



48 INORGANIC CHEMISTRY 

« 

It should be noted that the reduction of gas volumes to 
standard conditions is a mathematical process, and does 
not imply that the gas itself must actually be subjected to 
these conditions. (See Laboratory Manual, Appendix B.) 

Weight of a Liter of Oxygen Gas. — As already stated, the 
weight of a liter of oxygen is 1.429 gm. at 0° C. and 760 mm. 
Its weight would be different at any other temperature and 
pressure (unless the effect of heat and pressure balanced 
each other). This value (1.429) is found by an experiment 
involving several steps, (a) Oxygen is generated from a 
mixture of potassium chlorate and manganese dioxide, and 
the quantity liberated is found by subtracting the weight 
of the oxygen generator after the experiment from its original 
weight; suppose the weight of liberated oxygen is 2.312 gm. 
(b) The oxygen is collected, its volume noted, and the tem- 
perature and pressure also read; suppose the volume of 
oxygen is 1.75 1., the temperature is 19° C, and the pressure 
is 755 mm. (c) The observed volume is reduced to the 
volume at standard conditions, thus : — 

1.75 x 755/760 x 273/292 = 1.62. 

(d) The weight of one liter at 0° C. and 760 mm. is then 
found to be 1.427 gm. by dividing 2.312 (the weight of the 
oxygen) by 1 .62 (the corrected volume of the oxygen), thus : — 

2.312 -h 1.62 = 1.427. 

Very accurate experimental work involving precautions not 
mentioned above yields the value 1.429 gm. as the exact 
weight of one liter of oxygen. 

Densities of Gases. — The weight of a liter of many gases 
can be found by a method similar to that used for oxygen. 
The values obtained are comparable because the volumes 
are corrected for temperature and pressure (as reduction 



SOME PROrERTIES OF GASES 49 

to standard conditions is sometimes designated). Com- 
parison of the weight of a liter of different gases reveals 
interesting relations. Thus, a liter of air under standard 
conditions weighs 1.293 gm. Now if we divide 1.429 by 
1.293, the quotient (1.105) shows that oxygen is 1.105 times 
heavier than air. If we divide 1.429 by .08987, which is 
the weight of a liter of hydrogen at 0° C. and 760 mm., 
the quotient (15.9) shows that oxygen is 15.9 times heavier 
than hydrogen. This number (15.9) is the density of 
oxygen on the hydrogen standard. Density is the relative 
weight of equal and comparable volumes. Sometimes 
density is defined as the weight of a body in grams divided 
by its volume in cubic centimeters. Thus, according to 
the latter definition the density of oxygen is 1.429 -s- 1000 = 
.001429. But since in actual practice we more often com- 
pare gas volumes with each other, especially with hydrogen 
or air, the former definition {i.e. relative weights of equal 
volumes) is more convenient. Important deductions are 
made from the densities of gases, as will appear in a sub- 
sequent chapter. 

Problems 

1. Reduce the following volumes to the volume occupied at 
0°C. : (a) 173 cc. at 120° C, (6) 466 cc. at 14° C, (c) 706 cc. at 15° C, 
(d) 25 cc. at 27° C. 

2. Reduce the following volumes to the volume occupied at 
760 mm.: (a) 200 cc. at 740 mm., (6) 25 cc. at 780 mm., (c) 467 cc. 
at 756 mm. Arts, (a) 194.7, (6) 25.65, (c) 464.54. 

3. Reduce the following to the volume at standard conditions: 
(a) 147 cc. at 570 mm. and 136.5° C, (b) 320 cc. at 950 mm. and 
91° C, (c) 480 cc. at 380 mm. and 68.25° C, (d) 25 cc. at 780 mm. 
and 27° C, (e) 14 cc. at 763 mm. and 11° C. 

4. (a) The temperature of a gas is 18° C. At what temperature 
would its volume be doubled? (b) A gas measures 195 cc. at 740 
mm. What is the pressure when the volume is 295 cc. ? 



CHAPTER V 



General Properties of Water 

Water is a compound of hydrogen and oxygen, and is 
worthy of extensive study because of its indispensable re- 
lation to life, characteristic properties, and numberless uses. 

Occurrence in Nature. — Water in the form of vapor is 
always present in the atmosphere. Evaporation is con- 
stantly taking place from the surface of the ocean and other 
bodies of water, from the moist earth, from the bodies of 
animals, and from plants. This vapor condenses, and appears 
as clouds, mist, fog, snow, hail, dew, and frost. 

In the liquid state water occurs in vast quantities. About 
three fourths of the surface of the globe is water. Soil and 
porous rocks hold considerable, and plants and animals con- 
tain a large proportion. Certain substances, which are 
apparently dry, really retain much water. Thus, in a ton 
of clover hay there are upwards of 200 lb. of water. Many 
common foods consist largely of water, as may be seen by 
the following — 

Table of the Proportion of Water in Food 



Food 


Per Cext 
of Water 


Food 


Per Cent 
of Water 


Cod 


82.6 


Tomatoes 


94.3 


Beef 


61.9 


Apples .... 






84.6 


Lobster 


79.2 


Strawberries . . 






90.4 


Eggs 


73.7 


Watermelon . . , 






92.4 


Asparagus 


94.0 


Milk 






87.0 


Potatoes 


78.3 


Cheese .... 






28 to 72 


Cucumbers 


95.4 


White Bread . . 






35.3 



50 



GENERAL PROPERTIES OF WATER 51 

The human body is nearly 70 per cent water, and during 
a year the average man drinks about half a ton. 

Water in the form 06 ice permanently covers the coldest 
parts of the surface of the earth, e.g. the polar regions and 
the summits of high mountains. 

Functions of Water in Nature. — Since water is the only 
liquid occurring in large quantities on the earth's surface, 
it is the most effective agent of erosion. It cuts aw^ay the 
earth's crust, and transports the material from higher to 
lower levels, or washes it ultimately into the ocean. Acting 
in conjunction with carbon dioxide gas, it decomposes the 
rocks, changing them into clay, sand, and substances which 
make the soil productive. The cycle of changes from liquid 
to vapor and vapor to liquid exerts a marked influence on 
the distribution of heat and moisture upon the earth's 
surface, i.e. on climate. 

Water dissolves many solids and gases, and is constantly 
extracting from the rocks and soil their soluble constituents, 
some of which serve for the nutrition of plants, though 
the larger part passes on to the ocean. The latter thus 
becomes a vast reservoir of water containing salt and other 
mineral matter obtained from the earth's crust. In the 
vital processes of animals and plants water helps change 
the food into a condition fit for distribution and assimilation. 

Industrial Applications. — Besides the universal use of 
water as a beverage, it is applied to an endless variety of 
useful and convenient purposes. It has always been man's 
beast of burden. It is the vehicle for transferring me- 
chanical energy to water wheels — an application now being 
made on a vast scale for generating electricity. It utilizes 
by its peculiar properties the energy in fuel by means of 
the steam engine. It is the highway for transportation on 
the largest scale by ocean, rivers, lake, and canal. It is 



52 INORGANIC CHEMISTRY 

the vehicle for the distribution of heat by hot water and 
steam. It is the indispensable solvent in metallurgy, in 
the manufacture of many chemicals, and in such industries 
as soap making, bleaching, brewing, dyeing, and tanning; 
it is necessary wherever mortar and cement are used. Man's 
work would be stopped in a thousand other ways were he 
deprived of water. 

Natural Waters. — Water is never found pure in nature. 
Even rain water, which is the purest natural water, con- 
tains gases and dust washed from the air. When rain 
strikes the ground, it begins at once to take up impurities 
from the rocks and soil. Some of the water flows along 
the surface, becoming more and more impure, and finally 
reaches the ocean. But 25 to 40 per cent of the annual 
rainfall in temperate regions soaks into the ground and 
percolates through the soil at an estimated rate of .2 to 
20 feet a day. On its journey underground the water loses 
most, often all, of its organic matter {i.e. vegetable or animal 
matter or products of their decomposition), but it dissolves 
mineral matter and gases. The mineral matter is usually 
common salt and compounds of calcium and magnesium ; 
the most common gas is carbon dioxide. If the amount 
of dissolved matter in spring w T ater is large, or the kind of 
matter is so unusual as to give the water a marked taste or 
medicinal properties, the water is called mineral water. 
Water containing calcium and magnesium compounds is 
hard, but in soft water, such as rain water, these compounds 
are absent. 

• There are several hundred mineral springs in the United States, 
Those having a high temperature are called thermal, as at Hot Springs, 
Arkansas. Many contain a large proportion of common salt, as at 
Saratoga, New York. Others contain alkaline substances and carbon 
dioxide gas, e.g. Vichy and Apollinaris water. Sulphur springs con- 
tain solid or gaseous compounds of sulphur, — or both, — and have 



GENERAL PROPERTIES OF WATER 



53 



valuable medicinal properties. Some, like Hunyadi, are bitter; but 
others, especially those in New York State, which contain gaseous 
sulphur compounds, have a sweet taste but an unpleasant odor. Cha- 
lybeate waters contain soluble iron compounds. Many mineral 
waters contain calcium and magnesium compounds, and a few con- 
tain alum and lithium compounds. Most natural mineral waters 
contain traces of a large number of different substances. Many com- 
mercial mineral waters have indifferent medicinal value. 

River water obviously contains the impurities brought 
by springs and the surface water; it is also often made very 
impure by decaying animal and vegetable matter, which 
has been purposely or accidentally introduced, especially 
if the river passes through a thickly settled region. A 
sluggish river is more apt to be impure than a swift one, 
because the latter tends to purify itself by exposing its 
impurities to the oxidizing power of the air. Ocean water 
contains a large proportion of common salt. The propor- 
tions of the solid substances in their order of abundance 
are shown in the following — 

Table of Solid Substances in the Ocean 



Substance 


Per Cent 


Substance 


Per Cent 


Sodium Chloride . . . 
Magnesium Chloride . . 
Magnesium Sulphate . . 
Calcium Sulphate . . . 


77.76 

10.88 

4.74 

3.60 


Potassium Sulphate . . 
Calcium Carhonate . . 
Magnesium Bromide . . 
Other Substances . . . 


2.46 
.34 
.22 

traces 



The peculiar taste of ocean water is due to the presence 
of these substances, and since the water only is removed 
by evaporation, the ocean always has a " salty " taste. 

Drinking Water. — Water used as a beverage should 
of course be as pure as possible. As a rule the mineral 
matter in water selected for drinking is not injurious to 
health; but since water may become contaminated with 



54 INORGANIC CHEMISTRY 

bacteria which produce diseases such as typhoid fever and 
cholera, it is usually necessary to purify the water before use. 

The problem of obtaining suitable drinking water in large 
quantities is local. The water of many cities is purified 
by filtering it through a layer of sand and gravel, an acre 
or more in area and several feet deep. Such a filter removes 
bacteria almost completely, though it must be frequently 
cleaned. Sometimes the water is stored in a large settling 
basin or reservoir and purified by adding alum, or a similar 
substance, which causes the suspended matter to settle. 
Ozone is used as a purifier in some localities, and bleaching 
powder has been applied with excellent results to stored 
water contaminated with certain kinds of organic matter. 
Dissolved substances cannot be removed without consid- 
erable difficulty, so as a rule water is taken from a source 
which is reasonably pure. 

The purity of drinking water is usually determined by 
a water analysis. This is not a decomposition of water, 
but a chemical examination of a sample for the presence 
and amount of certain substances which indicate or cause 
impurity. A chemical examination is of limited value, 
however, unless it is supplemented by a microscopic study 
of a fresh sample and a rigid sanitary inspection of the 
premises. Water which is clear, sparkling, cool, attractive 
to the eye, and pleasant to the taste may be seriously pol- 
luted by disease germs ; or it may be liable to sudden con- 
tamination from some unsuspected source. On the other 
hand, a rather unpleasant-looking water may be harmless; 
hence the necessity of careful and extended examination of 
water to be used as a beverage. 

The purification of water may be readily accomplished 
by distillation. This operation consists essentially in boil- 
ing the water, condensing the resulting vapor, and collecting 
the liquid ; by this method the non-volatile matter dissolved 



GENERAL PROPERTIES OF WATER 



55 



in the water remains behind in the distilling vessel. It is 
performed in the laboratory by means of a condenser, which 
is shown in Figure 7. The condenser consists of an outer 
tube, A A, provided with an inlet and an outlet for a current 
of cold water, which surrounds the inner tube, BB. The 
vapor from the water boiling in the flask, C, condenses 
in the inner tube, owing to the decrease in temperature, 




Fig. 7. — ■ Condenser arranged for the distillation of water. 

and drops off the lower end of this tube, as the distillate, 
into the receiver, D, while the non-volatile impurities remain 
behind in the flask. Distilled water is prepared on a large 
scale in metal vessels, and the vapor is condensed in a block 
tin pipe coiled inside a vessel through which cold water 
flows. Distilled water is used in the chemical laboratory; 
large quantities are made into ice, and considerable (espe- 
cially after aeration) is used as a beverage. Distillation is 
done on a large scale by boiling water in a metal vessel 
and condensing the vapor in a spiral tin tube cooled by 
water. 



56 INORGANIC CHEMISTRY 

Physical Properties of Pure Water. — Owing to its marked 
solvent power, water is never found pure in nature, and is 
purified even in the laboratory only by taking special pre- 
cautions. At the ordinary temperature pure water is a 
tasteless and odorless liquid. It is usually colorless and 
transparent, but thick layers are bluish. Water is a poor 
conductor of heat. This last property can be shown by 
boiling water near the surface in a large test tube containing 
a piece of ice weighted down upon the bottom. The ice 
remains unmelted for some time, although the water is 
boiling a few inches above it. 

Most liquids expand when heated and contract when 
cooled. Water behaves exceptionally. If water at 100° C. 
is gradually cooled, it contracts until 4° C. is reached; if 
the cooling continues, the water expands as long as the liquid 
state is maintained. Hence at 4° C. a given volume contains 
the greatest weight of water. In other words, water has 
its maximum density at 4° C. The density of water at 4° C. 
is taken as 1 ; and water at this temperature is the standard 
for determining the specific gravity of solids and liquids. 
Thus, when we say specific gravity of gold is 19, we mean 
that a piece of gold is 19 times heavier than an equal volume 
of water at 4° C. A cubic centimeter of water at 4° C. 
weighs 1 gm. 

The expansion of water when cooled from 4° C. to 0° C. 
is slight, but the change is exceedingly important in nature. 
When the water on the surface of a lake or river cools, it 
contracts, and since it is heavier (volume for volume) than 
the warmer water beneath, it sinks. The warmer water 
rises, becomes cool, and likewise sinks, thus causing a cir- 
culation which continues until all the water from the sur- 
face to the bottom has a temperature of 4° C. Now if the 
cooling continues, the surface water expands and remains 
on the top, because it is lighter than the water beneath. 



GENERAL PROPERTIES OF WATER 57 

Hence when the temperature of the air falls to 0° C, this 
upper layer of water freezes and protects the remaining 
water from the cold air, thus stopping the circulation. 
Should the circulation continue, as the temperature fell 
from 4° C. to 0° C, the whole body of water would finally 
freeze from top to bottom. This condition would not only 
destroy aquatic life, but profoundly affect climate. 

Water solidifies or freezes at 0° C. (or 32° Fahrenheit). 
And when water freezes, it expands about one tenth of its vol- 
ume. That is, 100 cc. of water produce about 110 cc. of ice. 
In other words, 100 cc. of water and 110 cc. of ice weigh 
100 gm. each. Hence ice floats, The specific gravity of ice is 
about .92. The pressure exerted by water when it freezes 
is powerful. Vessels or pipes filled with water often burst 
when the water freezes. It is an erroneous but popular 
idea that " thawing out " a pipe bursts it. As a matter 
of fact, ice contracts when it melts. The pipe cracks as 
soon as the water freezes, and when the ice melts a channel 
is left for the water to flow out of the pipe. Because of this 
property, ice is an effective agent in splitting rocks. Water 
creeps into the cracks, especially into narrow ones, by capil- 
lary attraction, and when it freezes, the rock splits. Water 
in freezing also destroys the tissue of living plants, which 
are often said to have been "touched by frost. " Frozen 
flesh for a similar reason becomes pulpy and is more liable 
to putrefy when thawed. 

Ice melts at 0° C. (32° F.), which is also the freezing point 
of water. Ice often crystallizes in freezing, but the individ- 
ual crystals are seldom visible except during the first stages 
of the process. Snow crystals are common. They are 
always six-sided, and are formed in the atmosphere by the 
freezing of water vapor. 

Water evaporates at all temperatures, passing off as an 
invisible vapor into the atmosphere or into the air confined 



58 



INORGANIC CHEMISTRY 



over it. If water is heated in an open vessel, the tem- 
perature rises and vapor passes off rapidly until the ther- 
mometer reaches 100° C. (or 212° F.). At this point water 
boils ; i.e. it changes rapidly into vapor without rise of tem- 
perature. This vapor, if allowed to escape, cools and con- 
denses quickly into a cloud of minute drops of water. This 
cloud is popularly called steam. Accu- 
rately speaking, steam is invisible. What 
we call steam is a cloud or collection 




G 



678 




Fig. 8. — Vapor pressure. 

of very small particles of water. This may be illustrated 
by boiling water in a large glass flask. The inside of the 
flask is perfectly transparent, although there is a cloud of 
" steam " issuing from its mouth. 

Escaping vapor exerts pressure, as may be readily shown. 
If a little water is introduced into a dry closed bottle having 
a U-shaped tube of colored liquid connected with the interior 
to serve as a gauge, the difference in the levels of the colored 
liquid indicates a pressure inside the bottle (Fig. 8). This 
pressure is due to vapor escaping from the water. The 
ability of water (or of any other volatile liquid) to generate 
vapor is called the vapor tension of the liquid. The pressure 
exerted by the vapor is called the vapor pressure of water. 



GENERAL PROPERTIES OF WATER 59 

The amount of vapor pressure depends on the temperature. 
This is seen by comparing the heights of the mercury in the 
barometer tubes in Figure 8 (right). In the tube A there is 
no water vapor in the space above the mercury, and there- 
fore the height of the mercury is 760 mm. In the tube B 
the space above the mercury is filled with water vapor at 20° 
C. ; the vapor exerts a pressure and forces the mercury down 
to nearly 742 mm. That is, the water vapor at 20° C. exerts 
a pressure equal to about 18 mm. of mercury. Similarly, in 
the tube C the space is filled with water vapor at 50° C. and 
the mercury is forced down to 678 mm., the water vapor 
exerting a pressure of about 82 mm. If the vapor were at 
0° C, the vapor pressure w T ould be about 4.5 mm., and at 
100° C. the vapor pressure would be 760 mm. The latter 
value is instructive, for it means that at the boiling point of 
water (100° C.) the vapor pressure just balances the normal 
atmospheric pressure. The pressure exerted by water vapor, 
as shown in Figure 8, is independent of atmospheric pressure, 
the size and shape of the inclosing space, and the pressure 
of other gases. It depends solely on the temperature of 
the evaporating water, and has a maximum value for each 
temperature. The maximum value represents the vapor 
tension of water at the given temperature (see Equilibrium, 
page 61). These values can be found in the App., § 4. 

A practical application of vapor pressure is made in de- 
termining the weight of a liter of oxygen and in similar 
experiments where gases are measured over water. The 
oxygen gas is collected in a bottle or graduated tube inverted 
in a vessel of water. If the gas is allowed to stand confined 
over the water long enough, it becomes saturated with water 
vapor; i.e. the tube finally contains a mixture of oxygen 
and the maximum amount of water vapor at the given tem- 
perature. In such a mixture, where no chemical action 
occurs, each gaseous constituent shares the total pressure 



60 INORGANIC CHEMISTRY 

(against the atmospheric pressure). This proportionate 
part of the total pressure is called the partial pressure of 
that gas. If the oxygen were confined over mercury, its 
pressure would be the same as its partial pressure when 
saturated with water vapor. Hence the actual pressure 
under which the oxygen itself exists is found by determining 
its partial pressure or by subtracting the partial pressure 
of the water vapor from the total pressure (indicated 
by the barometer). The latter method is used, because the 
pressure of water vapor at any temperature is known and 
can be taken directly from the table. Incorporating this 
fact into the formula given in Chapter IV for reducing 
the volume of a gas to its volume at 0° C. and 760 mm. the 
formula becomes — 

V (P' - a) 



V = 



760 [1 + (.00366 t)] 



In this formula V means the volume of dry oxygen at 0° C. 
and 760 mm., and a means the vapor pressure (found in the 
table in the App., § 4). (See also Laboratory Manual, App. B.) 
Several important conclusions can be drawn from the 
discussion of vapor pressure in the foregoing paragraphs. 
Recalling the fact that the vapor pressure of water at 100° C. 
is 760 mm., it is obvious that this temperature is the boiling 
point at this pressure. Water boils when its vapor escapes 
with sufficient pressure to overcome the pressure upon its 
surface. Hence the boiling point depends upon the pressure 
— either of the atmosphere or of the vapor in the vessel. 
The boiling point of water is 100° C. (or 212° F.) only when 
the atmospheric pressure is normal, i.e. 760 mm. The boil- 
ing point becomes lower as the pressure is decreased and 
higher as the pressure is increased. For example, in the 
city of Mexico (7500 ft. above the sea level) water boils at 
92° C, and in Quito (9350 ft. above the sea level) water 



GENERAL PROPERTIES OF WATER 61 

boils at about 90° C. Again, the condition in a closed 
vessel containing water and water vapor merits considera- 
tion. As the water evaporates, the vapor pressure increases 
until the space above the water becomes saturated with 
water vapor. Then the vapor pressure remains constant as 
as long as the temperature is fixed, say 20° C. The vapor 
tension of the water is now equal to the vapor pressure of the 
vapor, In other words, vapor tension and vapor pressure 
balance each other. Vapor is condensing, however, just as 
fast as water is changing into vapor. Such a condition of 
mutual exchange is called a state of equilibrium, i.e. a con- 
dition in which two opposing processes balance but do not 
stop each other. Equilibrium is approached or reached in 
many processes, chemical as well as physical. 

Chemical Properties of Water. — Water has such con- 
spicuous physical properties that its chemical properties 
are sometimes ignored or discussed in connection with other 
substances. Such discussion, it must be admitted, is often 
more appropriate elsewhere, but certain phases of this 
topic need attention here. Water at the ordinary tempera- 
ture interacts with certain metals, especially sodium and 
potassium. This chemical change, as already stated, is 
substitution. The metal is substituted for part of the 
hydrogen of the water, thereby liberating hydrogen and 
producing a new compound, which contains sodium, hydro- 
gen, and oxygen, and is called sodium h}^droxide. Mag- 
nesium and zinc interact similarly with boiling water, the 
products being hydrogen, magnesium hydroxide, and zinc 
hydroxide. Water is decomposed to some extent into its 
component elements (oxygen and hydrogen) by intense 
heat; at about 2000° C. the decomposition is less than 2 
per cent. As the temperature falls, the elements recombine 
to form water. Both decomposition and recombination 



62 INORGANIC CHEMISTRY 

are gradual. Water when decomposed by heat behaves 
differently from potassium chlorate. The latter compound 
suffers permanent decomposition; i.e. the oxygen and the 
potassium chloride into which it is decomposed do not re- 
combine as the temperature falls. It is customary to 
distinguish these two kinds of decomposition. That kind 
which takes place gradually at high temperatures and is 
followed by recombination when the temperature is lowered 
is called dissociation by heat. Several common compounds 
undergo dissociation when heated, and they will be discussed 
later. Water is decomposed when heated with certain 
metals, e.g. iron, but the oxygen at once combines with 
the metal, so the final products are hydrogen and an oxide 
of the metal. A mixture of water and sulphuric acid yields 
hydrogen and oxygen when subjected to the action of an 
electric current. Apparently the water is merely decom- 
posed into its elements, but it will be shown subsequently 
that the action is not so simple. (See Chapter IX.) Water 
combines directly with many oxides. Thus, lime, which 
has the chemical name calcium oxide, combines directly 
with water and forms a compound called calcium hydroxide. 
Similarly, barium oxide forms barium hydroxide, sulphur 
dioxide forms sulphurous acid, phosphorus pentoxide forms 
phosphoric acid. These chemical changes may be typically 
represented thus : — 

Calcium Oxide + Water = Calcium Hydroxide. 
Sulphur Dioxide + Water = Sulphurous Acid. 

Oxides which react thus with water are called anhydrides. 
(See page 162.) Water also combines with certain solids 
when they separate from a solution by crystallization. 
Thus, copper sulphate crystals are blue, but when heated, 
water is liberated and the crystals crumble to a gray white 
powder. This class of compounds, in which water is a 



GENERAL PROPERTIES OF WATER 63 

definite component, is treated below. (See in this chapter 
Solution and Crystallization.) 

Solvent Power of Water. — This is one of the most con- 
spicuous properties of water and can be discussed from 
several standpoints. Only the simpler physical aspects are 
treated in this chapter. More extended discussion may be 
found in Chapter IX. 

Daily experience shows that many solids, liquids, and 
gases disappear when put into water. This operation is 
called dissolving or putting into solution. The clear, 
transparent liquid containing the dissolved substance is 
called the solution. The liquid in which the substance 
dissolves is called the solvent, and the dissolved substance 
is called the solute. If the solute is not volatile, or not very 
volatile, it can be recovered by evaporating, or distilling off, 
the water. The degree of solubility of a substance is often 
conveniently expressed by the terms slightly soluble, soluble, 
and very soluble. It is more accurate, however, to state 
the proportions of the solvent and the solute, and also the 
temperature. Thus, instead of saying that common salt is 
very soluble in cold w T ater, it is better to state that 36 gm. 
of salt dissolve in 100 gm. of water at 20° C. There is a 
limit to the weight of each substance which a given weight 
of water will dissolve. That is, substances differ widely 
in the degree of solubility. Some, like potassium perman- 
ganate, are very soluble, while others, like sand, dissolve 
only very slightly. The latter class is often said to be in- 
soluble. Strictly speaking, no substance is insoluble, but 
the term is often applied to those substances whose solu- 
bility is so slight that it can be neglected in most cases. 
Many minerals and rocks belong to the class of so-called 
insoluble substances. 

The concentration of a solution is its strength as de- 



64 INORGANIC CHEMISTRY 

termined by the amount of solute dissolved in a given 
amount of solvent. The ratio of these two amounts ex- 
presses the concentration. A solution which contains a 
small proportion of solute is called a dilute solution, or one 
of small concentration; one containing a large proportion 
of solute is called a concentrated solution. Thus, dilute 
sulphuric acid usually contains one part of acid to three 
or more parts of water, while concentrated sulphuric acid 
is nearly 98 per cent acid. Sometimes the terms weak 
and strong are used instead of dilute and concentrated. 
Other terms are defined below. (See Solubility of Solids.) 

Solution of Gases. — Water dissolves many gases. The 
solubility varies widely. For example, one volume of water 
dissolves about 1150 volumes of ammonia gas (at 0° C. 
and 760 mm.), 550 of hydrochloric acid gas, 80 of sulphur 
dioxide gas, 1.8 of carbon dioxide gas, .04 of oxygen, .02 of 
nitrogen, and .02 of hydrogen. 

As a rule the volume of gas dissolved by a given volume 
of water decreases with rise of temperature. Thus, when 
ammonium hydroxide is heated, ammonia gas escapes 
freely. So also, when water is heated gradually, the air 
dissolved in the water gathers as little bubbles, which soon 
rise and escape. Again, 100 volumes of water dissolve 
approximately 4 volumes of oxygen at 0° C, 3 at 20° C, 
1.8 at 50° C, and none at 100° C. The complete removal 
of a gas by boiling its solution is not possible in the case 
of certain very soluble gases, like hydrochloric acid gas. 
Such solutions when heated lose gas or water until a certain 
concentration is reached, and then the solution boils at a 
constant temperature (see pages 206, 218). 

Pressure influences the solubility of gases. Thus, large 
quantities of carbon dioxide gas are forced into cylinders 
full of water in preparing soda water. .When the pressure 



GENERAL PROPERTIES OF WATER 65 

is removed, the gas escapes rapidly and causes the soda 
water to froth or foam. This rapid escape of gas is called 
effervescence. Underground waters often contain large 
amounts of gases, especially carbon dioxide, owing to the 
great pressure to which subterranean gases are subjected. 
Hence, mineral waters often effervesce when they come to 
the surface. The greater the pressure, the greater the 
amount of gas dissolved. More accurately stated, the weight 
of a moderately soluble gas dissolved by a given weight of 
water is directly proportional to the pressure, if the tem- 
perature is constant. This is one form of the general state- 
ment known as Henry's Law. Gases which are very soluble 
or which dissolve under great pressure deviate from this 
law, probably owing to chemical combination. (See also 
preceding paragraph.) 

When a mixture of gases, such as air, dissolves in water, 
each ingredient behaves independently of the other. More 
strictly, each gas dissolves proportionally to its partial 
pressure. For example, the ratio of oxygen to nitrogen in 
air is about 1 : 4 and the ratio of their solubilities when not 
mixed is 2:1, but the ratio of oxygen to nitrogen in water 
saturated with air (at 760 mm.) is 7:13 or about 1:2. The 
difference is due to the fact that their solubility when mixed 
is determined not merely by their power to dissolve as free 
gases but also by their partial pressure. 

Solutions of Liquids. — The solubility 'of liquids in water 
varies between wide limits. Some, such as alcohol and 
glycerin, are soluble in all proportions. Oils, such as 
kerosene, are practically insoluble; hence the old adage, 
.-" Oil and water will not mix." Carbon disulphide is also 
almost entirely insoluble, as is shown by the fact that after 
agitation with water it separates almost entirely as a distinct 
layer ; being heavier than water, this layer forms at the 



66 INORGANIC CHEMISTRY 

bottom. The mere formation in this way of separate layers 
by two liquids is not conclusive evidence of relative insolu- 
bility. Only in those cases which exhibit perfect mutual 
solubility is the separation into layers after agitation im- 
possible. Ether and water form two layers, but each dis- 
solves appreciably in the other. Thus, about 2 gm. of 
ether dissolves in 100 gm. of water (at 20° C.) and about 
10 gm. of water dissolves in 100 gm. of ether. The upper 
layer consists of ether saturated with water; the lower, 
of water saturated with ether. Alcohol and water form no 
such layers, not because each is soluble in the other, but 
because each is soluble without limit in the other ; i.e. it is 
a case of perfect mutual solubility. In many cases a rise in 
temperature increases the solubility of liquids in water. 

Solutions of Solids. — The solubility of solids in water is 
a subject of much practical importance. The abundance 
of water and its power to dissolve such a vast number of 
different solids have led some to call water " the universal 
solvent/' The far-reaching effect of this marvelous power 
in nature and the indispensable value of water to man have 
been considered. (See above.) 

The degree of solubility of solids in water depends upon 
the substance itself and the temperature of the water. Some 
solids are very soluble, while others are difficultly soluble. 
In most cases solubility increases with a rise of tempera- 
ture ; hence the common practice of heating to hasten solu- 
tion. The effect of an increase of temperature on solubility 
is sometimes very marked, the solubility being increased 
many fold in passing from the ordinary temperature to the 
boiling point. A few solids (e.g. calcium hydroxide) are less 
soluble in hot water than in cold, and a few others (e.g. sodium 
chloride) dissolve to about the same degree in hot and cold 
water. These properties of the solutions of various solids 
are illustrated by the following — 



GENERAL PROPERTIES OF WATER 67 

Table of the Solubility of Solids in Water 



Solids 


Number of Grams in Solution in 
100 Grams of Water 




20° C. 


100° c. 


Calcium Chloride 

Copper Sulphate (cryst.) 

Magnesium Sulphate 

Potassium Chlorate 

Potassium Chloride 

Potassium Dichromate 

Potassium Nitrate 

Potassium Sulphate 

Sodium Chloride 


74.5 
42.3 
36.2 

7.2 
35 
13 
31.6 
10.6 
36 


159 
203.3 

73.8 

55.9 

57 
102 
246 

26 

39.8 



Other facts about the solubility of solids are illustrated by 
this table besides the dependence on the substance itself 
and on temperature. Inspection of the table shows that 
there is a limit to the solubility of solids in water at a given 
temperature. That is, a given weight of water at a fixed 
temperature will dissolve a definite weight of solid and no 
more, even though some undissolved solid is in the liquid. 
A solution conforming to the conditions just stated is said 
to be saturated or to have reached its maximum concentra- 
tion. Thus, 100 gm. of water holds 7.2 gm. of potassium 
chlorate in solution at 20° C. If more potassium chlorate is 
added, it remains undissolved (provided the temperature of 
the solution and the weight of water are unchanged). If 
the temperature falls, more solid comes out of solution ; 
and conversely, if the temperature rises, more solid dissolves. 
As long as the maximum concentration is maintained in 
contact with some of the undissolved solid, the solution is 
saturated. The ratio of the weight of the solute to the weight 
of the solvent in a saturated solution is called the solubility 
of the solid. Solubility is expressed in several ways. One 



68 



INORGANIC CHEMISTRY 



way represents the solvent by 100 gm. ; on this basis the 
solubility of a solid becomes the number of grams of solid 
dissolved by 100 gm. of water. In the case just cited, the 
maximum concentration, or, as it is more often called, the 




10° 20 



30 



40° 50° 60° 

Temperature 
Fig. 9. — Solubility curves. 



solubility, of potassium chlorate is 7.2 gm. at 20° C. and 
55.9 gm. at 100° C. 

The table of solubilities just given is limited to two tem- 
peratures. A convenient way of showing the solubility of 



GENERAL PROPERTIES OF WATER 69 

a substance as the temperature varies between convenient 
points is by a solubility curve. The curves of several sub- 
stances are shown in Figure 9. The temperatures are read 
from the vertical lines and the number of grams of solute in 
100 gm. of water is read from the horizontal lines. For 
example, if we wish to know the temperature at which 
40 gm. of potassium chlorate are held in solution by 100 gm. 
of water, it is only necessary to find where the horizontal 
line numbered 40 cuts the potassium chlorate curve, and 
then follow the vertical line down to the temperature num- 
bers, where 80° C. is found. Similarly, 100 gm. of water 
dissolve 37 gm. of sodium chloride at 60° C, while the same 
weight of water dissolves 110 gm. of potassium nitrate at 
60° C, and so on. 

When hot solutions are cooled or concentrated solutions 
are evaporated, the solute separates in the solid state just 
as soon as the saturation point (at a lower temperature) is 
passed. Under certain favorable conditions the solid is 
deposited in masses having a definite form. These masses 
are called crystals and the process of obtaining them is 
called crystallization. The form and color of the crystals 
are characteristic of the particular substance and serve to 
identify it. Thus, potassium chlorate crystallizes in shin- 
ing white plates or leaves, arid common salt in w^hite cubes. 
The deposition of crystals is not always as prompt as just 
stated. Thus a hot, concentrated solution of some solids, 
such as sodium acetate and sodium thiosulphate, deposits no 
crystals even when the clear solution cools. Such solutions 
are called supersaturated. Supersaturation can occur only 
when the undissolved solid is not present. If a fragment of 
the solid is dropped into the supersaturated solution, crystals 
very soon begin to form upon the fragment, and this separa- 
tion continues until just enough solid is left in solution to 
produce saturation at the prevailing temperature. The 



70 INORGANIC CHEMISTRY 

amount of solid thus separated is often very great and 
sometimes forms a solid mass in the test tube. 

It is evident from the foregoing paragraphs that there are 
three general classes of solutions of solids, viz. unsaturated, 
saturated, and supersaturated. These classes can be dis- 
tinguished by bringing each in contact with more of the 
solid. If the solution is unsaturated, more solid will dis- 
solve; if saturated, no more will dissolve; if supersaturated, 
solid will be deposited until saturation is reached. In an 
unsaturated solution the concentration is less than in a 
saturated solution, while in a supersaturated solution it is 
greater. If we think of these solutions as being in contact 
with a solid, their relations will be clearer than if we regard 
them as reservoirs, so to speak, for more or less solid. In the 
saturated solution there is equilibrium between the solution 
and the solid. That is, there is an equal tendency for the 
solid to enter and to leave the solution. But in the other 
two solutions no such equilibrium prevails, for the un- 
saturated solution takes up more solid and the supersatu- 
rated solution deposits solid; both are stable when solid is 
absent, but unstable as soon as solid is present. 

Solution and Crystallization. — Under the chemical prop- 
erties of water it was stated that water combines with cer- 
tain solids when they are separated from a solution by 
crystallization. Crystals deposited from solutions often 
contain water, which is an essential part of the compound. 
The combined water must not be confused with water which 
adheres to a crystal or is inclosed in it. Even after the crys- 
tals are powdered and dried, the combined water remains, 
which can be removed by heat or sometimes merely by ex- 
posure to the air. Loss of water is sometimes attended by 
loss of color and always by loss of crystalline appearance. 
Thus, crystallized sodium carbonate turns dull and crumbles 



GENERAL PROPERTIES OF WATER 



71 



in the air; blue crystallized copper sulphate turns white 
slowly at the ordinary temperature and very rapidly when 
heated, finally becoming a gray powder; but the variety of 
crystallized gypsum called selenite must be heated before 
the combined water passes off, whereupon the crystal be- 
comes chalky and crumbles when compressed. The pro- 
portion of combined water in crystals is not arbitrary. It is 
constant in the same compound when crystallized under 
uniform conditions. The amount in different substances 
varies between wide limits, as can be seen by the following — 

Table of Combined Water in Crystals 



Crystallized Solid 


Per Cent of Combined Water 


Barium Chloride 

Copper Sulphate 

Iron Sulphate 


14.75 
36.36 
45.35 



There is no doubt that the water in the crystallized form 
of compounds is not chemically combined in the way the 
other elements are ; for when such crystals are heated with 
proper precautions, only the water is removed, the other 
constituents of the original compound remaining intact. 
In most cases gentle heating suffices to expel the water. 
Hence the combination between the water and the rest of 
the compound must be weaker than that between the con- 
stituents of the residue or at least of a different order. This 
fact is often indicated by separating the water in the formula ; 
e.g. CuS0 4 . 5 H 2 and BaCl 2 . 2 H 2 0. Water chemically com- 
bined in a crystal and readily removed in a definite propor- 
tion by heating is called water of crystallization. Compounds 
containing water of crystallization are sometimes called 
hydrates or hydrated compounds. Conversely, compounds 
which have been deprived of water of crystallization are said 



72 INORGANIC CHEMISTRY 

to be anhydrous or dehydrated. For example, blue crystal- 
lized copper sulphate is a hydrate of the compound copper 
sulphate CuS0 4 , but when the blue compound is heated, it 
becomes anhydrous or dehydrated copper sulphate, which 
is a gray powder. Anhydrous compounds often readily be- 
come hydrated again. Thus, when the gray anhydrous 
copper sulphate is added to water, a blue solution is pro- 
duced from which blue crystals of hydrated copper sulphate 
are readily obtained. 

Some substances give up their water of crystallization 
wholly or in part upon mere exposure to the air; such sub- 
stances are said to be efflorescent or to effloresce. Thus, 
sodium carbonate crystals when left in the air become 
opaque and ultimately crumble, owing to the slow escape 
of their water of crystallization. The white spots often seen 
on blue crystals of copper sulphate or on green ones of 
iron sulphate are due to efflorescence. An explanation of 
efflorescence is found in the principle of vapor pressure. 
Substances containing water of crystallization exert a vapor 
pressure. If this vapor pressure is greater than the pressure 
of the water vapor in the atmosphere, the substance loses 
water until the vapor pressures are equal or until all the 
water has escaped from the substance. Hence, in general, 
all hydrated compounds effloresce, if they exert a vapor 
pressure greater than that of the atmosphere (at that time). 

It should not be concluded from the foregoing statements 
that all crystallized solids which dissolve in water contain 
water of crystallization. Many do not ; e.g. sodium chloride, 
potassium chlorate, sugar, potassium nitrate, and potassium 
dichromate. No satisfactory explanation has been given 
for the absence of water of crystallization in certain crystal- 
lized compounds; nor for the varying amount and relation 
to color and crystal form in hydrated compounds. 

The formation of crystals is not limited to a single process, 



GENERAL PROPERTIES OF WATER 73 

viz. cooling or evaporating a solution. It can be accom- 
plished by fusion and by sublimation. Thus, sulphur 
crystallizes when melted (or fused) and then cooled, and 
iodine crystallizes when vaporized and then cooled. These 
processes of obtaining crystals are called respectively evapo- 
ration, fusion, and sublimation. All three will be illustrated 
in succeeding chapters. As a rule crystals obtained by any 
of these methods are quite pure, and crystallization is one 
of the processes frequently used in the industrial preparation 
of chemicals. A crystal is u a solid body bounded by plane 
surfaces arranged according to definite laws, and possessing 
definite physical properties; both external form and physical 
properties resulting from, and being the expression of, definite 
internal structure." The external form is the most con- 
spicuous characteristic of a crystal, and, as a rule, each sub- 
stance has a crystal form or series of closely related forms 
by which it can be distinguished. Thus, salt crystallizes In 
cubes, alrTrh in octahedrons, sulphur in orthorhombic forms, 
and calcite (calcium carbonate) in hexagonal forms. Many 
minerals occur as crystals, and since they are the natural 
form of chemical compounds as well as the source of many 
elements, a knowledge of the common crystal forms and 
the properties of crystals is indispensable in identifying sub- 
stances and interpreting descriptions. Such knowledge is 
best acquired by a constant examination of crystals. The 
preliminary treatment given in the Appendix, § 3, will serve 
as an introduction to this important subject. 

Solution and Vapor Pressure. — Aqueous solutions have 
a vapor pressure, but it is less than the vapor pressure of 
water at the same temperature. This fact can be illustrated 
by introducing a solution into the barometer tube as described 
under vapor pressure. Under parallel conditions the mer- 
cury will always stand higher in the tube containing the 



74 INORGANIC CHEMISTRY 

solution, i.e. the vapor pressure is less, and consequently the 
mercury will be less depressed. Several conclusions can be 
drawn from this fact. Since the boiling point depends upon 
pressure, solutions must be heated to a higher temperature 
than water before boiling occurs. That is, a solid dissolved 
in water elevates the boiling point of water. Moreover, the 
elevation of the boiling point depends in general upon the 
weight of solid dissolved in a given weight of water ; i.e. it 
is proportional to the concentration of the solution. This 
means that if the weight of solute is doubled, the elevation 
of the boiling point is doubled. Similar statements can be 
made about the freezing point of solutions. The freezing 
point of a solution is lower than the freezing point of water, 
and the depression of the freezing point is proportional to 
the concentration of the solution. Important deductions 
will subsequently be made from these relations. In pass- 
ing, it is interesting to note two common illustrations of the 
foregoing statements, viz. that boiler water containing much 
dissolved solid has to be heated to a higher temperature than 
pure water in the production of steam, and that the salt 
water along the seashore freezes, if at all, with more diffi- 
culty than the fresh water of near-by rivers. Finally, the 
relatively lower vapor pressure of solutions explains the de- 
liquescence of certain substances. Many substances absorb 
water when exposed to the air, become moist, and sometimes 
even dissolve in the absorbed water. Calcium chloride, 
potassium carbonate, zinc chloride, sodium hydroxide, 
magnesium chloride, and potassium hydroxide belong to 
this class. This property is called deliquescence, and the 
substances are said to deliquesce, or to be deliquescent. 
Deliquescence is a property of very soluble substances. 
Water vapor from the air condenses on the surface and pro- 
duces a very concentrated solution, which has a vapor 
pressure much lower than the average pressure of the water 



GENERAL PROPERTIES OF WATER 75 

vapor in the air. The solution, therefore, continues to take 
up water until its vapor pressure equals the partial pressure 
of the water vapor in the air. Common salt, or sodium 
chloride, often deliquesces, especially in damp weather. The 
deliquescence is due, however, to the presence of magnesium 
and calcium chlorides. Sodium nitrate is somewhat deli- 
quescent, and is not usually used in the manufacture of 
gunpowder; potassium nitrate, which is not deliquescent, is 
used instead. This property of deliquescence is often utilized 
in the laboratory to remove water vapor from gases, calcium 
chloride being usually employed for this purpose. Substances 
thus used are often called drying or desiccating agents. 

Thermal Phenomena of Solution. — The process of solu- 
tion is often accompanied by an appreciable change of tem- 
perature. Thus, when sulphuric acid is poured into water, 
heat is liberated. With relatively large quantities of acid 
the heat is so great that the mixture often boils, and some- 
times the hot acid is spattered. Hence, the acid should be 
added slowly to the water, and the mixture constantly stirred. 
Other substances which dissolve with the liberation of heat 
are fused calcium chloride, potassium hydroxide, and sodium 
hydroxide. Some substances which dissolve with a fall of 
temperature (i.e. with absorption of heat) are crystallized 
calcium chloride, ammonium nitrate, ammonium chloride, 
and potassium nitrate. The final result of the entire act of 
solution is doubtless due to several factors. One of these is 
the change in volume of the solute. Thus, a very soluble 
gas becomes greatly condensed, while a solid, on the other 
hand, occupies a much larger volume after solution than 
before. Another important factor is chemical action be- 
tween the solute and solvent. The heat liberated or ab- 
sorbed during the process of solution is called heat of solution. 
(See Chapter XL) 



76 INORGANIC CHEMISTRY 

Solution and Chemical Action. — When a substance dis- 
solves, it is so modified that it can participate more readily 
in chemical changes. Hence, solution is an aid to chemical 
change. Thus, if dry tartaric acid and sodium bicarbonate 
are mixed, there is no evidence of chemical action; but when 
the mixture is poured into water, the copious evolution of 
carbon dioxide gas is conclusive evidence of a chemical change. 
Similarly, w T hen a dry mixture of ferrous sulphate and potas- 
sium ferrocyanide is poured into water, the immediate 
appearance of a blue precipitate shows that water was 
needed for the chemical change. Solution is such an im- 
portant aid to chemical action that many substances em- 
ployed in the laboratory are in solution, and many processes 
in chemistry are "wet" processes. (See Chapter IX.) 

The Nature of Solution has long been a subject of specula- 
tion and study. The problem as a whole is still unsolved, 
though much light has been thrown upon the question by 
recent investigations. (See Chapter IX.) 

Problems 

1. If 1.5 gm. of crystallized barium chloride loses .22 gm. when 
heated to constant weight, what per cent of water of crystallization 
does it contain ? 

2. If 2 gm. of another lot of barium chloride loses .295 gm., 
what per cent of it was water of crystallization ? 

3. A tube contains 97.2 cc. of gas at 20.3° C. and 756 mm., and the 
vapor pressure is 17.65. What is the volume of the dry gas at 0° C. 
and 760 mm. ? 

4. Reduce the following as in Problem 3: (a) 77 cc, 17.5° C, 
755 mm., 14.89 a; (b) 81.2 cc, 746.8 mm., 19.5° C, 16.87 a; 
(c) 100 cc, 755.3 mm., 18.5° C, 15.85 a. 

5. If the density of ice is .92, what volume will a liter of water 
at 4° C. occupy when frozen? Arts. 1.087 1. 

6. How much water (approximately) is contained in (a) 2 lb. 



GENERAL PROPERTIES OF WATER 77 

of lobster, (b) 56 lb. of potatoes, (c) 1 lb. of tomatoes, (d) 2 lb. of 
milk, (e) 1 lb. of white bread, (/) a human body weighing 150 lb. ? 

7. If a dry vessel of 1000 cc. capacity has a drop of water put 
into it at 20° C, what weight of water will evaporate? (Note. — 
One liter of water vapor at 0° C. and 760 mm. weighs .8045 gm.) 

8. 50 cc. of dry hydrogen at 18.3° C. and 758.7 mm. would 
occupy what volume at 0° C. and 760 mm. when saturated with 
water vapor ? 

9. Plot the following data on cross-section paper and draw 
the solubility curve of the substance : Temperature — 0, 10, 20, 
40, 55, 80 ; corresponding solubility (i.e. grams soluble in 100 gm. 
of water) — .8, .946, 1.18, 1.7, 2.1, 3.1. 

10. As in Problem 9 : Temperature — 0, 10, 20, 30, 40, 50, 60, 
70, 80, 90, 100; solubility — 26.9, 31.5, 36.2, 40.9, 45.6, 50.3, 55, 
59.6, 64.2, 68.9, 73.8. 

11. 30 cc. of a solution weigh 33.315 gm. and give 8.865 gm. of 
solid on evaporation. Calculate (a) the solubility of the solid, and 
(6) the weight of the solid in 100 cc. of the solution. 

12. If a cake of ice weighs 280 kg., what is its volume? 

13. A solution has a specific gravity of 1.8. How many cubic 
centimeters of water must be added to a liter of it to reduce its 
specific gravity to 1.5? 

14. By use of the solubility curves on page 68 answer : (a) How 
many grams of potassium chloride are in solution at 10°, 20°, 25°, 
60°, 80°, 95°, 100°? (b) Compare the solubility of potassium, 
chlorate, potassium bromide, potassium chloride, and potassium 
nitrate. How much of each is in solution at 30°, 50°, 90° ? (c) As 
in (b) — sodium chloride, ammonium chloride, lead nitrate, and 
sodium nitrate. 



CHAPTER VI 
Composition of Water — Hydrogen Dioxide 

Water was called an element until about the end of the 
eighteenth century. At that time it was shown to be a 
compound of hydrogen and oxygen. Since water is the first 
chemical compound we are to study, special attention will 
be paid to its typical characteristics. 

We should recall at this point the essential character- 
istic of a chemical compound, viz. its constituents are 
elements chemically combined in a fixed ratio by weight. 

The Composition of a Compound is determined either by 
analysis or synthesis, i.e. by taking it apart or putting its 
parts together. Sometimes both methods are used, since 
each fortifies the other and strengthens the final conclusion. 
These methods find excellent application in determining 
the composition of water. Analysis and synthesis may be 
qualitative or quantitative. A qualitative experiment is a 
study of compounds with a .view of discovering what ele- 
ments or groups of elements they contain. A quantitative 
experiment is an accurate determination of the weight or 
volume of the constituents of a compound. Obviously, a 
complete study of the composition of a compound requires 
both kinds of tests, the qualitative as a rule preceding the 
quantitative. 

Water contains Hydrogen. — When steam is passed over 
heated metals, hydrogen is liberated. Lavoisier's demon- 
stration of this fact has already been considered (see Prepara- 
tion of Hydrogen). The fact that sodium liberates hydrogen 

78 



COMPOSITION OF WATER 79 

from water at the ordinary temperature has also been dis- 
cussed (see ibid.). If red litmus paper is put into the water 
from which the sodium has liberated hydrogen, the litmus 
paper becomes blue. This change of color from red to blue 
shows that an alkali is in the water, because alkalies turn 
red litmus paper blue. The alkali is sodium hydroxide, and 
it may be obtained as a white solid by evaporating the water. 
Sodium hydroxide is a compound of sodium, hydrogen, and 
oxygen, and is formed by replacing half of the hydrogen of 
water by sodium. Since sodium liberates hydrogen from 
water, and forms at the same time a compound — sodium 
hydroxide — containing hydrogen, the hydrogen in water 
must be divisible into two parts. Now if .1 gm. of sodium 
is allowed to act upon water, 48.22 cc. of hydrogen are 
liberated; and if the sodium hydroxide thus formed is dried 
and heated with sodium, 48.22 cc. more of hydrogen are 
obtained. This shows that the hydrogen in water is divisible 
into two equal parts — a fact of fundamental importance. 

Water contains Oxygen. — The fact that oxygen is a 
constituent of water has already been suggested; e.g. (1) by 
the production of water when hydrogen is burned in air, 
(2) by the formation of a compound of iron and oxygen 
when steam is passed over hot iron, and (3) by the formation 
of sodium hydroxide when sodium interacts with water. 
These proofs, however, are indirect. A simple direct demon- 
stration of the presence of oxygen in water can be made 
by allowing chlorine water to stand in the sunlight. (Chlorine 
water is prepared by saturating water with chlorine gas — 
an element to be studied in Chapter XII.) A tube about a 
meter long is completely filled with chlorine water, the open 
end is immersed in a vessel containing the same solution 
and the whole apparatus is placed in the direct sunlight. 
Bubbles of gas soon appear in the liquid, and after a few 



80 



INORGANIC CHEMISTRY 



hours a small volume of gas collects at the top of the tube. 
This gas can be shown to be oxygen by the usual test, viz. 
relighting a glowing splinter of wood. 



Decomposition of Water. — When steam is heated to a 
very high temperature (about 2500° C), it decomposes to 

a very slight extent into hydro- 
gen and oxygen. This method, 
however, is not as convenient as 
the one in which an electric cur- 
rent is used. The decomposi- 
tion of water by electricity is 
called traditionally the elec- 
trolysis of water, though we 
shall see later (Chapters IX and 
XI) that the process is more 
complex than the apparent dis- 
ruption of water into hydrogen 
and oxygen. The operation is 
accomplished in a special form 
of apparatus devised by Hof- 
mann and shown in Figure 10. 
Pure water does not conduct 
electricity, so a mixture of water 
(10 vols.) and sulphuric acid 
(1 vol.) is poured into the ap- 
paratus until the reservoir is 
half full after the stopcocks have been closed. As soon as 
an electric battery of three or more cells is connected by 
wires with the piece of platinum near the bottom of each 
tube, bubbles of gas gather on the platinum, and as 
the action proceeds, the bubbles rise, collect in the upper 
part of the tubes, and slowly force the liquid from each 
tube into the reservoir. The volume of gas is greater in one 




Fig. 10. — Special form of Hof- 
mann apparatus for the elec- 
trolysis of an acid solution of 
water. 



COMPOSITION OF WATER 



81 



tube. Assuming that the tubes have the same diameter, the 
volumes are in the same ratio as their heights, which will be 
found by measurement to be approximately two to one. 
Appropriate tests show that the gas having the larger volume 
is hydrogen and the one having the smaller volume is oxygen. 
Many accurate repetitions of this experiment have shown 
that only hydrogen and oxygen are produced, and that the 
volume of the hydrogen is twice that of the oxygen. 

Water was first decomposed by electricity in 1800 by Nicholson 
and Carlisle. Davy confirmed their work by a series of brilliant 
experiments extending through a period of six years (1800-1806). 

The Quantitative Composition of Water. — Decisive evi- 
dence of the quantitative composition of water is obtained 
by a determination of its 
volumetric and its gravi- 
metric composition. Volu- 
metric means " by volume " 
and gravimetric means 
"by weight." 

The Volumetric Compo- 
sition of Water is deter- 
mined by exploding a 
mixture of known volumes 
of hydrogen and oxygen 
in a eudiometer. It is a 
method of synthesis. 

A simple sketch of a 
convenient form of appa- 
ratus for determining the 
volumetric composition of 
water is shown in Figure 
11. The essential part is the eudiometer, F. In this 
glass tube the gases are accurately measured and exploded. 




Fig. 11. — Apparatus for determining the 
volumetric composition of water. 



82 INORGANIC CHEMISTRY 

The electric spark which causes the explosion is obtained 
from ah induction coil and battery. The spark leaps across 
the space between the platinum wires at the top of the 
inside of the eudiometer, and the heat produced by this 
spark causes the hydrogen and oxygen to combine and 
form water. Omitting details, oxygen and hydrogen are 
introduced separately into the eudiometer and measured 
and the mixture is then exploded; after the explosion, which 
is indicated by a slight click or flash of light, water from the 
reservoir, E, rushes up into the eudiometer. The water 
does not completely fill the eudiometer, because an excess 
of one gas was added. This additional gas takes no part in 
the chemical change, but merely serves to lessen the violence 
of the explosion, which otherwise might break the eudiom- 
eter. The quantity of water formed by the union of the 
hydrogen and oxygen is too minute to measure. A concrete 
illustration will make the process more intelligible. Sup- 
pose the volumes of the participating gases after reduction 
to standard temperature and pressure and the dry state are 
as follows : — 

Vol. of hydrogen added 32.4 cc. 

Vol. of oxygen added 12.3 cc. 

Vol. of residual hydrogen . . . .. . 7.8 cc. 

The actual volume of hydrogen which combined with all 
the oxygen is 24.6 cc. (i.e. 32.4 — 7.8). Therefore the two 
gases combined in the ratio of 24.6 to 12. 3 ; or 2 to 1; or, as 
it is usually stated, two volumes of hydrogen combine with 
one volume of oxygen to form water. The ratio based on 
the most painstaking work is given as 2.0027 to_l. 

The discovery of the volumetric composition of water was not 
made by any one chemist. Priestley, about 1780, noticed that when 
a mixture of air and hydrogen was exploded, " the inside of the glass, 
though clear and dry before, immediately became dewy." Cavendish, 



COMPOSITION OF WATER 



83 



in 1781, showed that when a mixture of two parts hydrogen and one 
part oxygen was exploded, nothing but water was formed. Watt, in 
1783, was the first to state that water is a compound, though he per- 
formed no experiments and probably did not understand the real 
nature of its constituents. Lavoisier in the same year verified many 
facts previously noticed but not completely understood, and un- 
doubtedly first clearly recognized and stated what his contempo- 
raries had overlooked. The final proof of the volumetric composi- 
tion of water was an accurate verification in 1805 by Gay-Lussae 
and Humboldt of the previous observation that two volumes of 
hydrogen unite with one volume of oxygen. 

The Gravimetric Composition of Water is determined by 
passing dry hydrogen over copper oxide. The method de- 
pends upon the fact, previously stated, that many oxides, 
such as those of lead, copper, and iron, when heated in a 
current of hydrogen, give up their oxygen, or, chemically 
speaking, these oxides are reduced to metals. By this 
reduction the oxygen of the oxide combines with the hydro- 
gen, thereby forming water, which is collected in a weighed 
tube, while the metal remains behind in the original tube. 

A sketch of the apparatus is shown in Figure 12. The cop- 
per oxide is placed in the combustion tube, CC, which is 



^t^ir 




Fig. 12. — Apparatus for determining the gravimetric composition of water. 

made of hard glass. The Marchand tube, D, which is filled 
with calcium chloride, collects and retains the water formed 
in the combustion tube by the reduction of the copper 



84 INORGANIC CHEMISTRY 

oxide. The tubes A, B, and E keep moisture out of the 
apparatus. The experiment is easily conducted. Copper 
oxide is placed in the combustion tube, which is then care- 
fully weighed. The Marchand tube, filled with calcium 
chloride, is also weighed. After the other tubes are properly 
filled and are connected as shown in the figure, the hydrogen 
generator is adjusted so that a slow current can be passed 
through the whole apparatus (from left to right). The 
combustion tube is now heated, and moisture collects in it; 
as the heat increases, the copper oxide glows, and the mois- 
ture passes into the Marchand tube. When the operation 
is over and the apparatus is cool and free from hydrogen, the 
combustion tube and Marchand tube are weighed. The gain 
in weight of the Marchand tube is the weight of the water 
formed, while the loss in weight of the combustion tube is 
the weight of the oxygen removed from the copper oxide 
and now combined with the hydrogen in the form of water. 
An example will make this clear. Dumas, who first did 
this experiment accurately, found substantially that the 
combustion tube lost 5.251 gm. of oxygen, while the Mar- 
chand tube absorbed 5.909 gm. of water. Now the 5.909 
gm. of water contains .658 gm. of hydrogen (i.e. 5.909 — 
5.251). But .658 and 5.251 are in the same ratio as 1 and 
7.98. That is, water contains 1 part of hydrogen and 7.98 
parts of oxygen by weight. This ratio is very nearly 1 to 8, 
and the gravimetric composition of water is often stated as 
being 1 part hydrogen and 8 parts oxygen. Occasionally 
the gravimetric composition is stated in per cent, the values 
being 11.18 per cent hydrogen and 88.82 per cent oxygen. 
The ratio obtained by Dumas so long ago was scarcely ques- 
tioned until recently. We now know from an exceptionally 
accurate determination by Morley that the ratio is 1 to 
7.9395. He effected a complete synthesis of water in which 
the oxygen, hydrogen, and water were weighed. This ratio 



COMPOSITION OF WATER 



85 



ft 



an 



is now accepted as the correct one, though the more usual 
form is 2 to 15.879. The apparatus used by Morley is shown 
in Figure 13. It wag first weighed vacuous (i.e. free from 
air or other gases). The tubes, aa, were 
then connected with the weighed reservoirs jp a S . 

of oxygen and hydrogen, and the oxygen 
was introduced. Sparks were next passed 
between the platinum wires, cc, and the 
heat ignited the hydrogen, which was slowly 
admitted, the combination of the gases 
taking place at bb. The water vapor 
condensed in the tube dd, the lower por- 
tion of which was immersed in water. 
The combustion of the hydrogen was con- 
tinued until a suitable weight of water was 
formed. The water and its vapor were 
then converted into ice by putting the 
apparatus into a freezing mixture; the 
residual mixture of gases was drawn off 
and analyzed, passing in its exit through 
tubes of phosphorus pentoxide in ee which 
retained all traces of water. The whole 
apparatus was finally weighed, the increase 
being the weight of the water formed by the combination 
of known weights of hydrogen and oxygen. 

Summary. — Water is a chemical compound of hydrogen 
and oxygen combined in a fixed ratio by weight, viz. 1 to 
7.9395; they are also combined in the ratio of 2.0027 to 1 
by volume. Usually these ratios are stated approximately 
as 2 to 16 by weight and 2 to 1 by volume. 



Fig. 13. — Morley's 
apparatus for de- 
termining the grav- 
imetric composi- 
tion of water. 



The gravimetric composition of water was first determined about 
1820 by Berzelius and Dulong. Their work was verified by Dumas 
about 1842. The complete synthesis was made in 1895 by Morley. 



INORGANIC CHEMISTRY 

Problems and Exercises 

1. What (a) weight and (b) volume of hydrogen are needed 
to change 75 1. of oxygen into water? (Assume standard condi- 
tions.) 

2. What (a) weight and (b) volume of oxygen are needed to 
change 185 cc. of hydrogen into water ? (Assume standard conditions. ) 

3. Suppose 150 1. of water are decomposed by electricity. 
What (a) weight and (b) volume (at 0° C. and 760 mm.) of the 
products are formed ? 

4. How many (a) grams and (b) cubic centimeters (at 0° C. 
and 760 mm.) can be prepared from a metric ton of water? 

5. Apply Problem 4 to oxygen. What would be the volume 
of oxygen at 18° C. and 767 mm. ? 

6. If 20 gm. of water were produced by the explosion of a mix- 
ture of hydrogen and oxygen, what volumes of the dry gases were 
used at 15° C. and 770 mm. ? 

7. How many grams of potassium chlorate are needed to pro- 
duce enough oxygen for the complete combustion of 10 gm. of 
hydrogen ? 

8. What volume of (a) hydrogen and (b) oxygen, both at 12° C. 
and 762 mm., can be obtained by decomposing 10 gm. of water 
by electricity? 

9. Morley found that 3.2645 gm. of hydrogen combined with 
25.9176 gm. of oxygen. What is the gravimetric composition of 
water according to this experiment ? 

10. How would you prove that water is composed of only 
hydrogen and oxygen? 

11. State the sources of error in determining the gravimetric 
composition of water by the copper oxide method. 

12. State (a) the exact gravimetric and (b) the exact volumetric 
composition of water. Describe the method of determining each. 

Hydrogen Dioxide 

Hydrogen Dioxide is a liquid composed of hydrogen and 
oxygen. But the proportion of the constituents is not the 
same as in water. It contains approximately one part of 
hydrogen and sixteen parts of oxygen by weight. It is 
often called hydrogen peroxide, because its relative pro- 



COMPOSITION OF WATER 87 

portion of oxygen is greater than in water — the other 
hydrogen oxide. 

Preparation. — It is prepared by treating barium dioxide 
(or peroxide) with sulphuric acid. The commercial solu- 
tion has a variable strength, and usually contains 3 per cent 
or more of hydrogen dioxide. 

Properties. — Hydrogen dioxide has a sharp, pungent 
odor and a bitter, metallic taste. The concentrated solu- 
tion is a syrupy liquid, but the commercial solution is scarcely 
distinguishable in appearance from water. It is an unstable 
compound, which decomposes slowly at the ordinary tem- 
perature, and very rapidly if heated. The dilute, commercial 
solution is somewhat stable, but heat and light decompose 
it completely into water and oxygen. The ease with w T hich 
it yields oxygen makes it a good oxidizing agent. Under 
certain conditions it is also a reducing agent. 

Uses. — Dilute solutions are used extensively to bleach 
animal and vegetable matter, such as human hair, ostrich 
feathers, fur, silk, wool, cotton, bone, and ivory. It is also 
used as an antiseptic and disinfectant in surgery on account 
of its oxidizing properties. Large quantities are used to 
restore the color to faded paintings — a use suggested by 
Thenard, the discoverer. In the laboratory it is a service- 
able reagent. 



CHAPTER VII 

Law, Theory, and Hypothesis — Laws of Definite and Multiple 
Proportions — Atomic Theory — Atoms and Molecules — 
Symbols and Formulas — Equations 

Law and Theory. — We discover facts by observation and 
experiment. Phenomena which always occur under the 
same conditions soon become well-established facts. Our 
knowledge of the facts which have some relation to each 
other is often summarized in a brief statement called a law. 
A law is not only an epitome of the uniform behavior of 
observed facts. It is also a statement which permits us to 
predict occurrences under like conditions ; for if a law is valid 
for many observed cases, we conclude that it will cover future 
cases. The essential feature of a law is universal validity. 

The ultimate cause of scientific facts is unknown. The 
explanation we give of facts, especially of groups of related 
facts, is called a theory. Laws are statements about facts; 
theories are statements of the supposed cause of facts. 
Laws seldom change, but theories are often modified. Laws 
are the result of experiment, but theories are the outcome of 
mental operations. We accept a certain theory until a more 
satisfactory one is proposed. Sometimes experiment yields 
results which need further examination or are beyond the 
realm of our present experimental skill. The temporary 
explanation or supposition we make in such a case as a guide 
in further investigation is called an hypothesis. An hypothe- 
sis is often the forerunner of a theory, and they are not always 
sharply differentiated. 

Laws, theories, and hypotheses are of great service in 

88 



LAW, THEORY, AND HYPOTHESIS 89 

chemistry, since they permit us to gather into intelligible 
statements our knowledge of a vast number of related facts, 
as well as assist us in discovering new facts and interpreting 
the phenomena of nature. In this chapter we shall discuss 
two laws and one theory that are of great importance in 
chemistry. 

Law of Definite Proportions by Weight. — When the metal 
magnesium is heated in the air, it burns with a dazzling flame 
and yields a grayish powder, due to combination with oxy- 
gen. If magnesium is heated so that the product cannot 
escape, a remarkable relation between their weights is re- 
vealed. In order to burn completely 1.52 gm. of magnesium, 
1 gm. of oxygen is necessary; and the product, magnesium 
oxide, weighs 2.52 gm. This product contains, therefore, 
60.317 per cent of magnesium and 39.682 per cent of oxygen. 
Accurate repetitions of this experiment have shown that this 
proportion by weight is fixed and definite. Again, if a weighed 
quantity of potassium chlorate is decomposed, 31.903 per 
cent of potassium, 28.932 per cent of chlorine, and 39.164 
per cent of oxygen are always obtained. This means that 
the proportion of potassium, chlorine, and oxygen which 
makes up potassium chlorate is fixed and definite. Experi- 
ments similar to these show that in all of the chemical 
compounds which have been examined the different con- 
stituents are always present in a definite and unvarying 
proportion by weight. There are no exceptions to this 
general fact. This constancy of proportion in chemical com- 
pounds is stated as the Law of Definite Proportions by 
Weight, thus: — 

A given chemical compound always contains the same ele- 
ments in the same proportion by weight. 

This law is one of the fundamental laws of chemistry. It 
is so confidently believed that if the composition of a com- 



90 INORGANIC CHEMISTRY 

pound is found by analysis to vary, chemists conclude that 
the experimental work is incorrect or that the compound is 
impure. The law was established as the outcome of a con- 
troversy between two French chemists, Proust (1755-1826) 
and Berthollet (1748-1822). The discussion lasted from 
1799 to 1806. Berthollet believed that compounds might 
have a varying composition. Indeed, by his experiments 
he detected " gradual changes" in composition. But Proust 
showed that Berthollet analyzed mixtures and not com- 
pounds. (In a mixture the components may be present 
in any proportion.) Subsequent experiments have only 
strengthened our confidence in this law. 

Law of Multiple Proportions. — Proust showed that some 
elements combine in more than one proportion by weight, 
and thereby produce two or more distinct compounds. But 
he failed to notice that if the weight of one element is adopted 
as constant, the varying weights of the other element or ele- 
ments are in a simple multiple relation to each other. Dalton 
discovered this general fact. about 1804. The composition 
of compounds is usually expressed in per cent; but such 
expressions in a series of compounds reveal nothing about 
multiple relations. If, however, a constant weight of one 
constituent is adopted as a basis of comparison, and the com- 
position of the series of compounds is expressed in terms of 
this weight, then the simple multiple relation which exists 
between the weights of the other constituent or constituents 
is clearly seen. For example, no multiple relation is evident 
from the statement that two compounds contain respec- 
tively 27.27 and 42.8571 per cent of carbon and 72.72 and 
57.1428 per cent of oxygen. But if in expressing the com- 
position of these compounds we adopt some convenient 
number, such as 1 for the weight of carbon in each com- 
pound, the weights of oxygen will be in the simple integral 



LAW, THEORY, AND HYPOTHESIS 



91 



ratio of 2 to 1. The five compounds of oxygen and nitrogen 
which will soon be studied, aptly illustrate this fact of mul- 
tiple proportions : — 

Table to illustrate Multiple Proportions 



Name 


Composition in 
Per Cent 


Adopted 
Weight 


Ratio 




Nitrogen 


Oxygen 


Nitrogen 


Nitrogen 


Oxygen 


Nitrous Oxide . . . 
Nitric Oxide . . . 
Nitrogen Trioxide 
Nitrogen Peroxide . 
Nitrogen Pentoxide . 


63.636 

36.842 
30.434 
25.925 


36.363 
53.333 
63.157 
69.565 
74.074 


1 
1 
1 
1 

1 


1 

1 
1 
1 
1 


.57 
1.14 
1.71 

2.28 
2.85 



From this table it is clear that the weights of oxygen com- 
bined with the same weight of nitrogen are as 1 : 2 : 3 : 4 : 5 ; 
i.e. they are in a relation to each other w T hich can be expressed 
by small integral numbers. The same simple ratio would be 
obtained if any other value were substituted for 1, and similar 
tables may be worked out with other series of compounds. 

The general fact of multiple proportions is stated as tho 
Law of Multiple Proportions, thus: — 

When two or more elements, a., b, c, etc., unite to form a series 
of compounds, a fixed weight of a always combines with such 
weights of b (as well as of c, etc.), that the ratio between these 
different weights can be expressed by small (usually), whole 
numbers. 

This law, like the law of definite proportions, is a funda- 
mental law. And together with the law of the conservation 
of matter they have profoundly influenced the theoretical 
and practical progress of chemistry. 

The Atomic Theory. — The laws just discussed state in 
condensed form certain general facts about the quantitative 



92 INORGANIC CHEMISTRY 

aspects of chemical change. They point to the existence of 
chemical units which participate in chemical changes with' 
out alteration of weight. But we have no means of detect- 
ing or separating these chemical units. Nevertheless, in order 
to provide a mental picture of this quantitative feature of 
chemical change a theory has been proposed. It is called 
the atomic theory and was announced in approximately its 
present form by Dalton, an English chemist, about 1805. 
According to this theory, (1) chemical elements and com- 
pounds consist ultimately of a vast number of very small 
particles called atoms; (2) chemical change is union, separa- 
tion, or exchange of undivided atoms ; (3) atoms of the same 
chemical element are alike and have an unvarying weight 
called the atomic weight; (4) atoms of different elements 
differ from each other in weight. The atomic theory means 
in a few words that matter consists of atoms which are en- 
dowed with a weight characteristic of each element and which 
remain undivided in chemical changes. 

This theory, it will be observed, deals with the nature of 
matter and with the quantitative aspects of chemical change. 
It does not state facts nor does it make facts more valid. 
It serves merely to assign some explanation, more or less 
detailed, to facts and laws already formulated. Let us 
apply the atomic theory to certain facts, i.e. restate these 
facts in te'rms of the theory. Before proceeding with this 
application, however, it will be necessary to make a pre- 
liminary distinction between an atom and a molecule. About 
the time of Dalton (1766-1844) the term "particle was used to 
include both atom and molecule. But they are not iden- 
tical. An atom is the smallest particle of an element that 
participates in a chemical change. Molecules are particles 
which consist of two or more atoms chemically combined ; 
atoms are alike in molecules of elements but different in mole- 
cules of compounds. Thus, the smallest particle of copper 



LAW, THEORY, AND HYPOTHESIS 93 

which participates in chemical changes is an atom, and the 
smallest particle of water is a molecule, which consists of 
hydrogen and oxygen atoms chemically combined. A fuller 
discussion of atoms will soon be given. But this preliminary 
distinction will permit us to resume the application of the 
atomic theory to certain facts. An appreciable mass of the 
element copper consists of many millions of atoms of copper, 
all alike — all having the same unvarying weight. Sodium, 
oxygen, hydrogen, sulphur, carbon, and all the other elements 
likewise consist of atoms, but the weight of the sodium atom 
differs from the weight of the atom of copper, carbon, sulphur, 
and all the other elements. Again, when a chemical change 
occurs between copper and sulphur, for example, atoms of 
copper combine with atoms of sulphur and produce molecules 
of a compound called copper sulphide. And this combining 
of atoms into molecules continues until all the copper atoms 
or the sulphur atoms, or under special conditions the atoms 
of both substances, have been used. Furthermore, this 
chemical change takes place not only between vast numbers 
of atoms, but the quantitative aspects of this multitude of 
changes conform to the atomic theory. This latter point 
needs explanation, because it emphasizes the chief feature 
of the atomic theory, viz. agreement with certain fundamental 
laws of chemical change. These laws are the law of the con- 
servation of matter, the law of definite proportions, and the 
law of multiple proportions. (1) According to the atomic 
theory the weights of atoms do not change ; all other prop- 
erties may be temporarily lost or buried, but the weight is re- 
tained throughout all chemical changes however complex. It 
is obvious that if atoms never change their weights, the total 
weight of matter in a chemical change is unvaried, and the law 
of the conservation of matter follows as a natural consequence. 
(2) Again, according to the atomic theory, when magnesium 
combines with oxygen, molecules of magnesium oxide are 



94 INORGANIC CHEMISTRY 

formed by the union of some whole number of atoms of mag- 
nesium with some whole number of atoms of oxygen. Each 
molecule of magnesium oxide would therefore consist of one 
or more atoms of magnesium united with one or more atoms 
of oxygen, and the composition of each molecule of mag- 
nesium oxide would be definite; i.e. each molecule would 
contain the same elements united in a definite proportion by 
weight. In other words, magnesium oxide would always be 
found to consist of a certain per cent of magnesium and a 
certain per cent of oxygen. Hence the atomic theory har- 
monizes with the law of definite proportions. (3) Finally, the 
atomic theory conforms to the law of multiple proportions. 
The number of atoms of the combining elements may be 
the same or different. That is, the ratio of combination 
may be 1 to 1, 1 to 2, 2 to 3, etc. According to the atomic 
theory atoms are transferred as wholes; i.e. in chemical 
changes there are no fractions of atoms. Therefore, the 
proportions of the weights of different elements in a series of 
compounds must be simple proportions; all multiple rela- 
tions will be expressible by whole numbers. An illustration 
will make this point clear. There are two compounds of 
carbon and oxygen. Analysis shows that in the one contain- 
ing the less oxygen the ratio of the weight of carbon to oxy- 
gen is 3 to 4, and in the other 3 to 8. The second compound 
contains twice as much oxygen as the first; i.e. its molecule 
contains twice as many atoms of oxygen as a molecule of the 
first. (Subsequently it will be shown that the first com- 
pound is carbon monoxide and contains one atom of oxygen 
in each molecule, while the second is carbon dioxide and 
contains two atoms of oxygen in each molecule.) In other 
words, the weights of oxygen in this series of compounds are 
in the simple ratio of 1 to 2. The same line of reasoning can 
be applied to other series of compounds whatever the ratio 
of combination. Hence the atomic theory harmonizes with 
the law of multiple proportions. 



LAW, THEORY, AND HYPOTHESIS 95 

Atoms and Molecules. — It should not be overlooked that 
the laws of definite and multiple proportions deal with facts, 
while the atomic theory deals with conceptions which may 
be true, but which we cannot prove to be true. We often 
speak of atoms as if they could be perceived by the senses, 
but we do so simply because such methods of expression help 
us describe, study, and interpret chemical changes. We 
also describe elements as if they consisted of free atoms, but 
atoms do not, as a rule, exist in the uncombined state. As 
soon as atoms are freed from combination, they at once unite 
with some other atom or atoms. The particles which make 
up oxygen gas are not atoms, but a group or chemical com- 
bination of atoms. Groups of chemically combined atoms 
are called molecules. The molecule of an element consists 
of atoms of one kind only, but the molecule of a compound 
consists of atoms of two or more different kinds. For 
example, a molecule of oxygen consists of atoms of the ele- 
ment oxygen, while a molecule of the compound water con- 
sists of atoms of oxygen and hydrogen. (A molecule of a few 
elements contains only one atom.) Atoms are the chemical 
constituents of molecules. They are the smallest particles 
of the elements which are known to participate in chemical 
changes. 

Our views regarding molecules are based on extensive in- 
vestigation of the properties of gases. The molecule is often 
spoken of as the physical unit, because in most physical 
changes molecules are not decomposed. Whereas the atom 
is called the chemical unit because it is the part of a molecule 
which as a rule is transferred unchanged in chemical changes. 
Molecules will be discussed again. (See Chapter XIV.) 

Although the atom is conceived to pass as a whole from 
compound to compound, it should not be inferred that 
atoms cannot be decomposed under any conditions. The 
phenomena exhibited by compounds of radium show that 



96 INORGANIC CHEMISTRY 

there are particles smaller than atoms. (See Radio-activity.) 
These very small particles are called corpuscles or electrons. 
However, the atom is the chemical unit, and whether a 
single individual or a group of smaller individuals, its weight 
is not altered in chemical changes. 

Atomic Weights. — The essential property of matter is 
weight. According to the atomic theory, different kinds of 
atoms have different weights. But the absolute weight 
of an atom cannot be directly determined by any instru- 
ments available. We can, however, find the relative weight 
of an atom ; that is, how many times heavier one atom is than 
another atom. These relative weights are called the atomic 
weights of the elements. If we should adopt 1 as the weight 
of an atom of hydrogen, the weights of atoms of other ele- 
ments could be readily expressed in terms of this standard. 
Thus, the atomic weight of sodium would be 22.88, of oxy- 
gen 15.88, of carbon 11.9, etc.; that is, an atom of sodium 
weighs twenty-three times as much as an atom of hydrogen, 
etc. For many years hydrogen (=1) was the standard. 
But for scientific reasons oxygen is now the standard, and 
16 is adopted as its atomic weight instead of 15.88. This 
change does not alter any facts ; it merely changes slightly 
the numerical values of the atomic weights. Their relation 
to each other is not changed. The atomic weight of hydro- 
gen becomes 1.008, if oxygen equals 16, and others are pro- 
portionally changed. The atomic weights are real weights 
because they are found by experiment. It should be con- 
stantly borne in mind, however, that they are relative 
weights. That is, the atomic weight of sodium, for example, 
is 23.00, not 23.00 gm. or any other absolute weight, but 
23.00, if the atomic weight of oxygen is 16. 

The exact determination of the atomic weight of an ele- 
ment is a difficult operation. Many principles influence the 



LAW, THEORY, AND HYPOTHESIS 97 

final selection of the number to be adopted as the atomic 
weight. This subject is discussed in Chapter XIV. A 
complete table of the atomic weights is given in the Appendix, 
§ 5. Exact and approximate values of the atomic weights 
of the important elements can be found in a table on the 
inside of the back cover, and for most purposes these approx- 
imate values can be used, e.g. in solving the problems in this 
and subsequent chapters. 

Chemical Symbols, which were mentioned in Chapter I, 
represent single atoms. Thus, H represents one atom of 
hydrogen, O one atom of oxygen, N one atom of nitrogen. 
If more than one uncombined atom is to be designated, the, 
proper numeral is placed before the symbol, thus : — 

2 H means 2 uncombined atoms of hydrogen, 

3 O means 3 uncombined atoms of oxygen, 

4 P means 4 uncombined atoms of phosphorus. 

But if we wish to represent the atoms as in chemical com- 
bination, either with themselves or with other atoms, then a 
subscript is used instead of a coefficient, thus : — 

H 2 means 2 atoms of hydrogen in combination, 
N 3 means 3 atoms of nitrogen in combination, 
P 4 means 4 atoms of phosphorus in combination. 

Symbols not only represent atoms,, but they also express 
atomic weights. Thus, represents one atom of oxygen, but 
it also means that this atom weighs 16. Similarly, K rep- 
resents one atom of potassium, which weighs 39.10. 

Chemical Formulas. — A formula is a group of symbols 
which is designed to express the composition of a compound. 
A given compound, as we have already seen, has a definite 
composition. In other words, a. molecule of a given com- 



98 INORGANIC CHEMISTRY 

pound always contains the same number and the same kind 
of atoms ; and since the molecules are alike, the composition 
of one molecule is the same as the composition of the com- 
pound. In writing a formula, the symbols of the different 
atoms making up a molecule of the compound are placed 
side by side. Thus H 2 is the formula of water, because one 
molecule of this compound consists of 2 atoms of hydrogen 
and 1 atom of oxygen. Similarly, KC10 3 is the formula of 
potassium chlorate. These symbols might be written in a 
different order, but usage has determined the order in most 
cases. A formula, as just stated, represents one molecule. 
Hence, KC10 3 represents one molecule of potassium chlorate, 
and means that the molecule of this compound contains 1 
atom each of potassium and chlorine and 3 atoms of oxygen. 
If we wish to designate several molecules, the proper nu- 
meral is placed before the formula, thus : — 

2 KC10 3 means 2 molecules of potassium chlorate, 

3 H 2 means 3 molecules of water, 

4 H 2 S0 4 means 4 molecules of sulphuric acid. 

In certain compounds some of the atoms act like a single 
atom in chemical changes. This fact is often expressed by 
inclosing the group of atoms in a parenthesis, or by separat- 
ing it from the rest of the formula by a period. Thus, the 
formula of ammonium nitrate is (NH 4 )N0 3 , and the formula 
of alcohol is C2H5.OH. Water of crystallization is usually 
indicated in the same manner, CuS0 4 .5 H 2 being the for- 
mula of crystallized copper sulphate. The period is some- 
times omitted, especially if the composition of the compound 
is well understood. If a group of atoms is to be multiplied, 
it is placed within a parenthesis. Thus, the formula of lead 
nitrate is Pb(N0 3 ) 2 . This means that the group N0 3 is to 
be multiplied by 2. That is, lead nitrate might be expressed 
by PbN 2 6 , but for reasons which will be given later the 



LAW, THEORY, AND HYPOTHESIS 99 

formula Pb(N0 3 ) 2 is used. The expression 2 Pb(N0 3 ) 2 means 
that the formula Pb(N0 3 ) 2 is to be multiplied by 2. 

Symbols and formulas are sometimes used to represent an in« 
definite amount of an element or compound. Thus, O is often used 
to designate oxygen and H 2 S0 4 sulphuric acid. They are often thus 
used to label bottles in a laboratory. Such a departure from accuracy 
should not be allowed to obscure their real meaning. 

The complete significance of symbols and formulas can be grasped 
only by their intelligent use. They should not be committed to 
memory slavishly. It is desirable, however, to learn the common 
ones while the substances they represent are being studied, and 
leave the consideration of their relations until the needed facts have 
accumulated. 

Molecular Weights. —Since atoms combine to form mole- 
cules, a molecular weight is the sum of the weights of the 
atoms in the molecule. A molecule of nitric acid contains 

1 atom each of hydrogen and nitrogen, and 3 atoms of oxy- 
gen; hence its molecular weight is 1 + 14 .01 + (16 X 3) =63.01. 
Given the formula, the molecular weight is easily found by 
adding the atomic weights. Just as a symbol stands for an 
atomic weight, so a formula expresses a molecular weight. 
Molecular weights are real numbers, but they are not found 
experimentally by merely adding the atomic weights. For 
convenience we may add the atomic weights, but historically 
and experimentally molecular weights preceded exact atomic 
weights. Many facts and principles determine the selection 
of the molecular weight of a compound. These are discussed 
in Chapter XIV. For the present, empirical knowledge is 
sufficient. That is, it will answer all purposes to find the 
molecular weight of a compound by adding the atomic 
weights corresponding to the number of atoms in the formula 
of a molecule of the given compound. By this somewhat 
arbitrary procedure the molecular weighted: water (H 2 0) is 

2 + 16 = 18; the weight of 2 H 2 is 2(2 + 16) = 36. Simi- 
larly, the molecular weight of lead nitrate (Pb(N0 3 ) 2 ) is 207.20 



100 INORGANIC CHEMISTRY 

+ 2(14.01 + 48) = 331.22; the weight of two molecules of 
lead nitrate (2 Pb(N0 3 ) 2 ) is 2 X 331.22 = 662.44. Also, the 
molecular weight of sulphuric acid (H 2 S0 4 ) is (2 X 1) + 32.06 
+ (4 X 16) = 98.06. Other molecular weights may be calcu- 
lated in the same way. 

It should be noted that the molecular weight of a com- 
pound, like the atomic weight of an element, is a relative 
weight. That is, the molecular weight of water is not 18 
gm., but 18, if the atomic weight of oxygen is 16. 

Calculations based on Formulas and Molecular Weights. — 

It is evident from the foregoing paragraphs that there is a 
rigid connection between the molecular weight and the 
formula of a compound. Since the formula expresses the 
composition of a compound by means of small integral num- 
bers representing the ratio of the atomic weights in a mole- 
cule, it is possible to calculate (1) the composition in per cent, 
if the formula is known, and (2) the formula, if the compo- 
sition in per cent is known. Composition in per cent, or, as 
it is usually designated, percentage composition, is readily 
calculated from the formula of a compound. Let us take 
a concrete case. The formula of sulphuric acid is H 2 S0 4 . 
This formula represents a molecular weight of 98, i.e. 
2 + 32 + 64 = 98 (using approximate atomic weights). Now 
if the respective parts of hydrogen, sulphur, and oxygen (viz. 
2, 32, 64) are divided by 98 and the quotient then multiplied 
by 100 {e.g. -fa X 100), the product is the per cent of each 
element in sulphuric acid. It is sometimes more convenient 
to solve the problem by a proportion. Thus, the proportions 
for the percentage composition of sulphuric acid are : — 

2 : 98 : : X : 100; X = 2.04 per cent of hydrogen. 
32 : 98 : : X : 100; X = 32.65 per cent of sulphur. 
64 : 98 : : X : 100; X = 65.31 per cent of oxygen. 

Total, 100.00 per cent. 



LAW, THEORY, AND HYPOTHESIS 101 

By the same method the percentage composition of any com- 
pound can be calculated. The calculation of a formula, 
when the atomic weights and the percentage composition 
are known, is practically the converse of the above process; 
i.e. it is simply the process of finding the small integral 
numbers which are in the same ratio as the numbers express- 
ing the composition. Suppose we know the composition 
of sulphuric acid to be hydrogen = 2.04 per cent, sulphur 
32.65 per cent, oxygen 65.31 per cent. If the percentage of 
each element is divided by the corresponding atomic weight, 
the quotients are 2.04, 1.02, and 4.08. Reducing these quo- 
tients to integral numbers (by dividing by 1.02), the final 
quotients are 2, 1, 4. But these quotients represent the ratio 
of the atomic weights in a molecule; that is, the relative 
number of atoms of each element in a molecule. Therefore 
the formula of sulphuric acid must be H 2 S0 4 . The formula 
of a compound calculated by this method is its simplest 
formula. (See also Determination of Formulas of Com- 
pounds, Chapter XIV.) ^ 

Chemical Equations. — When substances interact chem- 
ically, the definite chemical transformation is called a 
reaction. We have already seen that reactions can be ex- 
pressed by equations. Thus, in Chapter II two of the reac- 
tions involved in the preparation of oxygen were expressed 
as follows: — 

Barium Oxide + Oxygen = Barium Dioxide. 

Barium Dioxide = Barium Oxide + Oxygen. 

It was stated at that point that these equations are crude 
forms of chemical equations. We can remodel these prelim- 
inary equations by using symbols and formulas in place of 
words, the equations just given then becoming : — 

BaO + O = Ba0 2 . 
Ba0 2 =BaO + 0. 



102 INORGANIC CHEMISTRY 

The preceding remodeled equations are ordinary chemical 
equations; i.e. they are equations showing the kinds and 
relative weights of the interacting substances. 

The scope and interpretation of the ordinary chemical 
equation can best be set forth by a further discussion of the 
four kinds of chemical changes. 

(1) Decomposition. When mercuric oxide is heated, it 
changes into mercury and oxygen. This reaction is ex- 
pressed by the equation : - — 

HgO = Hg + 

Mercuric Mercury Oxygen 
Oxide 

This equation may be read in several ways : (a) Mercuric 
oxide decomposes into mercury and oxygen; (b) one mole- 
cule of mercuric oxide by decomposition forms one atom of 
mercury and one atom of oxygen; (c) 216 parts by weight 
of mercuric oxide yield 200 parts by weight of mercury and 
16 parts by weight of oxygen (since these are the relative 
weights found by experiment and reduced to the atomic 
weight basis); (d) mercuric oxide equals mercury plus 
oxygen. 

(2) Combination. When magnesium burns in air or in 
oxygen, magnesium oxide is formed. The equation for the 
reaction is : — 

Mg + O = MgO 

Magnesium Oxygen Magnesium 
Oxide 

This equation may be read as follows : (a) Magnesium and 
oxygen combine to form magnesium oxide ; (b) one atom of 
magnesium combines with one atom of oxygen and forms 
one molecule of magnesium oxide; (c) 24 parts by weight 
of magnesium combine with 16 parts by weight of oxygen 
and yield 40 parts by weight of magnesium oxide; (d) mag- 
nesium plus oxygen equals magnesium oxide. 



LAW, THEORY, AND HYPOTHESIS 103 

(3) Substitution. When zinc and sulphuric acid interact, 
zinc replaces the hydrogen of the sulphuric acid. The equa- 
tion for this reaction is : — 

Zn + H 2 S0 4 = 2 H + ZnS0 4 

Zinc Sulphuric Hydrogen Zinc 

Acid Sulphate 

This equation may be read as follows : (a) Zinc and sulphuric 
acid interact and form hydrogen and zinc sulphate by the 
substitution of zinc for the hydrogen of the acid; (6) one 
atom of zinc interacts with one molecule of sulphuric acid 
and forms two atoms of hydrogen and one molecule of zinc 
sulphate; (c) 65 parts by weight of zinc interact with 98 
parts by weight of sulphuric acid and yield 2 parts by 
weight of hydrogen and 161 parts by weight of zinc sulphate; 
(d) zinc and sulphuric acid equal hydrogen and zinc 
sulphate. 

(4) Double Decomposition. When solutions of silver ni- 
trate and sodium chloride are mixed, silver chloride and 
sodium nitrate are formed. This kind of chemical change is 
called double decomposition, or metathesis, because both of 
the original compounds undergo decomposition. It is really 
an exchange or redistribution of atoms; the original com- 
pounds undergo decomposition while the final compounds 
result from a combination of these parts on another plan. 
Double decomposition may be regarded as the simultaneous 
occurrence of the other kinds of chemical change, for it in- 
volves decomposition, combination, and substitution. How- 
ever, it is customary to give this kind of chemical change 
a special name, owing to certain unique features displayed 
by it. One of these features is the quite frequent forma- 
tion of an insoluble solid called a precipitate. For example, 
in the reaction just cited the silver chloride is produced as 
a white, curdy solid, almost insoluble in the final liquid. 
Double decomposition often results in precipitation. It is 



104 INORGANIC CHEMISTRY 

therefore an excellent way to make tests, and it finds numer- 
ous applications in chemical analysis. The equation for the 
foregoing reaction is : — 

AgN0 3 + NaCl = AgCl +NaN0 3 

Silver Sodium Silver Sodium 

Nitrate Chloride Chloride Nitrate 

This equation may be read as follows : (a) Silver nitrate and 
sodium chloride interact and form silver chloride and sodium 
nitrate by double decomposition; (b) one molecule of silver 
nitrate interacts with one molecule of sodium chloride and 
forms one molecule of silver chloride and one molecule of 
sodium nitrate; (c) 170 parts by weight of silver nitrate 
interact with 58.5 parts by weight of sodium chloride and 
yield 143.5 parts by weight of silver chloride and 85 parts by 
weight of sodium nitrate; (d) silver nitrate plus sodium 
chloride equals silver chloride plus sodium nitrate. 

In the foregoing discussion of the four kinds of chemical 
change similar statements are designated by the same letter. 
Let us consider each lettered group. Under (a) in each case 
nothing is said about the physical conditions attending the 
chemical change recorded by the equation, because such 
accompaniments are outside the scope of ordinary chemical 
equations. Thus, in the equations HgO = Hg + O and 
Mg + O = MgO there is no hint whatever that the mercuric 
oxide must be kept at a high temperature or that the mag- 
nesium unites vigorously with oxygen at a relatively low 
temperature. Again, in the equations Zn + H 2 S0 4 = 2 H + 
ZnS0 4 and AgN0 3 + NaCl = AgCl + NaN0 3 there is no sug- 
gestion that the sulphuric acid, sodium chloride, and silver 
nitrate must be dissolved in water, nor that the zinc sulphate 
which is produced remains in solution while the silver chloride 
is precipitated. Furthermore, heat is often liberated in 
chemical changes, e.g. by the interaction of zinc and hydro- 
chloric acid; but in the ordinary chemical equation as given 



LAW, THEORY, AND HYPOTHESIS 105 

above this fact is ignored. Hence we conclude (1) that ordi- 
nary chemical equations tell nothing about the physical 
conditions (i.e. temperature, physical state, solution, etc.) 
under which the chemical reaction starts, proceeds, and 
ends. Again, in (b) and (c) the chemical equations are made 
up of the smallest integral number of atoms and molecules 
involved in the chemical change. The equation is a sample, 
so to speak, of the vast number of like changes which we 
call the chemical change. What determines the number of 
atoms and molecules to be incorporated in an equation? 
The answer to this question necessitates the discussion of 
several topics. First, equations are the outcome of ex- 
periments. They follow experiments, and are designed to 
be compact, symbolic expressions of certain phases of a par- 
ticular chemical change. They tell at a glance one part of 
a complex story. Second, ordinary chemical equations ex- 
press quantitative relations. That is, they not only empha- 
size the fact that a chemical change exemplifies the law of 
the conservation of matter, but they also show the propor- 
tions of the participating substances. For example, ex- 
periment shows that when a given weight of mercuric oxide 
is decomposed into mercury and oxygen, the actual weights 
involved are in the ratio of 216 to 200 to 16 respectively. 
Corresponding values are found for each equation. These 
values, which differ of course with different equations, are 
real and must be known before the particular equation can 
be written correctly. In fact, they precede every equation, 
although we often overlook this fact in using equations. 
Third, before a chemical equation can be written certain 
facts must be known. One of these is the composition of 
each compound involved; i.e. not merely the per cent of each 
constituent in each compound, but the proportion of each 
constituent in terms of the atomic weights of the elements. 
In other words, before an equation can be written it is 



106 INORGANIC CHEMISTRY 

necessary to know the symbol and atomic weight of each 
element and the formula and molecular weight of each com- 
pound. These, as previously stated, may be found for the 
present by utilizing the table of atomic weights given in the 
Appendix, § 5, supplemented by information given in the 
text. It is also essential to know the proportions in which 
the original substances (often called factors) interact and 
in which the final substances (sometimes called products) 
are produced. 

We draw as a second conclusion from this rather long dis- 
cussion of the statements recorded above (in (b) and (c)) that 
(2) an ordinary chemical equation shows by means of the 
appropriate number of atoms and molecules not only the 
kind but the relative quantities of substances before and 
after a chemical change. 

(d) The verbal interpretation of an ordinary chemical 
equation is very often compressed into a form which simu- 
lates the algebraic equation. But ordinary chemical equa- 
tions cannot be subjected to transposition or factoring. In 
a certain sense, therefore, it is anomalous to say that mercuric 
oxide equals mercury plus oxygen. There will be little or no 
difficulty, however, if it is understood that the ordinary chem- 
ical equation is not an algebraic expression which attempts 
to describe a chemical change in all its aspects. Objections 
have been raised to the use of the sign of equality (=), and 
an arrow is sometimes substituted for it, the equation then 
becoming, for example, HgO — > Hg + O. Many equations 
are used merely to express the gravimetric proportions in 
which chemical reactions take place ; in such equations we 
shall use the sign of equality. Doubtless certain facts can 
best be expressed by equations in which the sign of equality 
is replaced by an arrow. The special style or form of an 
equation depends upon the part of the chemical story it is 
designed to epitomize. Later we shall have occasion to use 



LAW, THEORY, AND HYPOTHESIS 107 

other forms of equations. (See ionic, thermal, and gas equa- 
tions, and equilibrium.) 

We draw finally a third and comprehensive conclusion ; 
viz. (3) ordinary chemical equations are expressions show- 
ing by symbols and formulas the quantitative relations be- 
tween all the substances involved in chemical reactions. 
Each equation is the outcome of experiment, and although 
the equation contains signs used in algebra, a chemical 
equation has none of the properties of an algebraic equation 
except equality between the total weights on each side of the 
equation ; the chemical equation, furthermore, is limited 
to a statement of the chemical distribution of atoms and 
does not reveal any facts about the physical phenomena 
which invariably accompany chemical changes. 

Making Equations. — It is clear that the task of making 
a correct chemical equation is not easy. Several avenues 
are open to the beginner. The short equations can be com- 
mitted to memory or worked out by methods soon to be 
outlined; the long ones can be interpreted by the facts 
recorded in connection with the experiment and then re- 
ferred to later, as occasion demands. The student must 
not forget that each chemical reaction has its own equation 
and that similarity of names and of chemical changes does 
not imply uniformity of equations. A word of caution 
must also be uttered against attempts to write an equation 
by mere guess work and then expect the facts to coincide 
with this pseudo-equation. 

One method of working out simple chemical equations 
will be clear from the following cases: (a) When magnesium 
is heated in oxygen (or in air), the ratio by weight in which 
' the two elements combine is 3:2. This result is expressed 
in terms of the atomic weights of the elements involved. 
Let y equal the number of atomic weights of magnesium 



108 INORGANIC CHEMISTRY 

and z the number of atomic weights of oxygen. Then we 
can write a preliminary equation thus : — 

y X at. wt. of magnesium : z X at. wt. of oxygen = 3:2. 

The atomic weight of magnesium is found by the table on 
the back cover to be 24 and that of oxygen to be 16. 
Substituting these values, we have the equation : — 

2/x24:2Xl6 = 3:2 

By inspection y = z, and the simplest value of each is 1. 
Now the symbol Mg expresses 24 parts of magnesium and 
the symbol O expresses 16 parts of oxygen. Therefore 
Mg and O are the symbols representing the smallest number 
of atoms equivalent arithmetically to the ratio (3 : 2) found 
by experiment. The formula of the product formed by 
their combination is therefore MgO, and the simplest equa- 
tion expressing the chemical change is — 

Mg + = MgO 

Magnesium Oxygen Magnesium 

Oxide 

(6) Again, suppose we wish to find the simplest equation 
for the reaction between hydrogen and oxygen in the for- 
mation of water. Experiment shows that hydrogen and 
oxygen combine in the ratio 1 : 8 by weight. Pursuing the 
same line of argument as above, let y = the number of atomic 
weights of hydrogen, and z that of oxygen. The preliminary 
equation is : — 

y X at. wt. of hydrogen : z X at. wt. of oxygen = 1:8 

The atomic weight of hydrogen is 1 and of oxygen is 16. 
The equation now becomes — 

y X 1 : z X 16 = 1 : 8 

By inspection, y = 2 z, and the simplest values are y = 2 
and z = 1. Now the symbol H stands for 1 part of hydrogen, 



LAW, THEORY, AND HYPOTHESIS 109 

and O for 16 parts of oxygen. Therefore, 2 H and are 
the symbols representing the smallest number of atoms 
equivalent arithmetically to the ratio (1:8) found by ex- 
periment. The formula of the product of their combina- 
tion is H 2 0, and the simplest equation is — 

2 H + = H 2 

All equations are not equally simple, but by a similar 
argument based upon the facts found by experiment many 
simple equations may be developed. 

Another method is often possible. When we know the 
factors and products of a reaction, we can find their symbols 
or formulas in the book, construct a preliminary equation, 
and then balance the equation; i.e. select the proper coeffi- 
cients, subscripts, or both, so that there shall be an equal 
number of atoms of each element on both sides of the equa- 
tion. An example will make this method clear. When 
phosphorus burns in oxygen, phosphorus pentoxide is formed. 
The preliminary equation is — 

P + = P 2 5 

Here it is evident that to balance the equation we need 2 P 
and 5 on the left. Hence the final equation is — 

2P + 50 = P 2 5 

Again, when zinc and hydrochloric acid interact, hydrogen 
and zinc chloride are formed. The preliminary equation 
made from the symbols and formulas is — 

Zn + HC1 = H + ZnCl 2 

By inspection, it is evident that two atoms of chlorine are 
on the right and only one on the left. To obtain Cl 2 it is 
necessary to write 2 HC1. But 2 HC1 means not only 2 CI 
but 2 H. Hence the equation becomes — 

Zn + 2 HC1 = 2 H + ZnCl 2 



110 INORGANIC CHEMISTRY 

A final inspection shows that an equal number of atoms of 
each element is on both sides of the equation. 

Many equations may be written by applying these methods 
to the facts found by experiment (see exercises at the end of 
this chapter). 

Equations for Preceding Reactions. — The equations cor- 
responding to many reactions already discussed may ap- 
propriately be collected here, partly for their retrospective 
value and partly for future use. The equation for the 
preparation of oxygen from mercuric oxide is — 

HgO = Hg + O 

Mercuric Mercury Oxygen 
Oxide 

When sulphur, carbon, magnesium, and iron are burned in 
air (or in oxygen), the equations are — 

S + 2 = S0 2 

Sulphur Oxygen Sulphur 
Dioxide 

C + 20 = C0 2 

Carbon Oxygen Carbon 
Dioxide 

Mg + O = MgO 

Magnesium Oxygen Magnesium 

Oxide 

3Fe+ 4 = Fe 3 4 

Iron Oxygen Magnetic 

Iron Oxide 

The equation for the preparation of hydrogen from zinc and 
hydrochloric acid is — 

Zn + 2 HC1 = 2 H + ZnCl 2 

Zinc Hydrochloric Hydrogen Zinc 

Acid Chloride 

When hydrogen burns or when a mixture of hydrogen and 
oxygen is exploded, the equation is — 



LAW, THEORY, AND HYPOTHESIS 111 

2 H + O = H 2 

Hydrogen Oxygen Water 

The equation for the reaction in determining the gravimetric 
composition of water is — 

CuO + 2H = H 2 + Cu 

Copper Hydrogen Water Copper 
Oxide 

The equation for the decomposition of potassium chlorate is — 
KCIO3 = KC1 +30 

Potassium Potassium Oxj^gen 
Chlorate Chloride 

The interaction of sodium and water is represented thus: — 
Na + H 2 = H + NaOH 

Sodium Water Hydrogen Sodium 

Hydroxide 

When phosphorus burns in air (or oxygen) , the equation is — 
2P + 50 = P 2 5 

Phosphorus Oxygen Phosphorus 
Pentoxide 

Calculations based on Equations. — Since equations are 
expressions of the relative quantities of the substances in- 
volved in chemical reactions, it is possible to solve many 
arithmetical problems arising from reactions. An equation 
states the proportions which participate chemically in a 
reaction. Obviously, any convenient weights of zinc and 
sulphuric acid might be brought together, but the pro- 
portions according to which- the factors react and the prod- 
ucts are formed are always expressed by the equation — 

Zn + H 2 S0 4 = 2 H + ZnS0 4 

Zinc 
65 

If zinc and sulphuric acid are brought together in any other 
proportion, a part of one or the other will be left over unused. 



Sulphuric 


Hydrogen 


Zinc 


Acid 




Sulphate 


98 


2 


161 



112 INORGANIC CHEMISTRY 

The equation above means that zinc and sulphuric acid 
always interact in the proportion of 65 and 98, and produce 
hydrogen and zinc sulphate in the proportion of 2 and 161. 
We may read grams, ounces, kilograms, or any other unit 
in connection with these numbers, but the same unit must 
be used throughout the calculations. Therefore, if we know 
the actual weight of one substance participating in a reaction, 
all other weights involved can be readily calculated. Sup- 
pose 45 gm. of zinc interact with sulphuric acid; the weights 
of (a) acid required, -(b) hydrogen formed, and (c) zinc 
sulphate produced are calculated as follows : — 

(1) Write the chemical equation for the reaction, thus: — 

Zn + H 2 S0 4 = 2 H + ZnS0 4 

(2) Place under each term of the equation its atomic or 
molecular weight, 1 as the case may be, thus: — 

Zn + H 2 S0 4 = 2H + ZnS0 4 

65 98 2 161 

(3) Place above the proper terms the known weight and 
required weight (i.e. X, Y, Z, etc.) involved in the problem, 
thus : — ■ 

45 X Y Z 

Zn + H 2 S0 4 = 2 H + ZnS0 4 

65 98 2 161 

(4) State in the form of a proportion the four terms in- 
volved, remembering that the known and required weights 
are in the same ratio as the atomic and molecular weights. 
Thus, the three proportions in the given problem are: — 

(a) 45 : X : : 65 : 98; X = 67.8 gm. sulphuric acid. 
(6) 45 : Y : : 65 : 2; Y = 1.38 gm. hydrogen, 
(c) 45 : Z : : 65 : 161 ; Z = 111.4 gm. zinc sulphate. 

Similar problems can be solved by this method. 

1 The atomic weights are given in the table on the inside of the back 
cover. Molecular weights are obtained by adding the proper atomic weights. 

/ 



LAW, THEORY, AND HYPOTHESIS 113 

Problems and Exercises 

1. Calculate the percentage composition of (a) water (H 2 0), 
(b) zinc oxide (ZnO), (c) lead carbonate (PbC0 3 ), (d) sodium 
chlorate (NaC10 3 ), (e) barium oxide (BaO), (/) calcium carbonate 
(CaC0 3 ). 

2. Calculate the percentage composition of (a) copper sul- 
phate (CuS0 4 ), (b) barium chloride (BaCI 2 ), (c) manganese dioxide, 
(d) calcium oxide, (e) sodium hydroxide, (/) potassium hydroxide, 
(g) sodium carbonate (Na 2 C0 3 ), (h) potassium nitrate (KN0 3 ), 
(i) mercuric oxide (HgO). 

3. Show that the following sets of compounds illustrate the 
law of multiple proportions : (a) H = 11.11 per cent and = 88.88 
per cent, H = 5.882 and = 94.117; (b) Sn= 62.63 and Cl = 
37.37, Sn = 45.49 and CI = 54.41. 

4. Calculate the formula of the compounds which have the 
indicated composition : (a) Na == 60.68, CI = 39.31 ; (b) Ca = 29.41 ; 
S = 23.52, = 47.05; (c) C = 27.27, = 72.72; (d) As = 75.8, = 
24.2; 0) N = 82.35, H= 17.63. 

5. Calculate the formula of the compounds which have the fol- 
lowing composition : (a) Si = 19.5, C = 66.62, H = 13.88; (6)Pb = 
86.6, S = 13.4 ; (c) N = 26.17, H = 7.48, CI = 66.35. 

6. How much oxygen can be prepared from (a) 70 gm. of mer- 
curic oxide; (b) 17 gm. of potassium chlorate? 

7. How much potassium chloride will remain after 82.5 1. of 
oxygen (at 0° C. and 760 mm.) have been obtained from potassium 
chlorate ? 

8. Interpret the following: H, H 2 , 2 H, H 2 0, 2 H 2 0, H 2 2 , 
NaOH, Ca(OH) 2 , A1(N0 3 ) 3 , ZnCl 2 , 2 ZnS0 4 , 3 Fe 3 4 , 5 P 2 5 . 

9. Interpret the following : (a) Pb0 2 = PbO + ; (b) Cu + = 
CuO; (c)Zn+2NaOH=Na 2 Zn0 2 + 2H; (d) Ba(N0 3 ) 2 + K 2 S0 4 = 
BaS0 4 + 2 KN0 3 . 

10. From the equation KC10 3 = KC1 + 3 0, calculate (a) the 
weight of potassium chloride when 30 gm. of oxygen are liberated ; 
(b) the weight of chlorine the potassium chloride will yield, if 10 
gm. of potassium chlorate are decomposed ; (c) the volume of 
oxygen liberated (at 0° C. and 760 mm.) from 35 gm. of potassium 
chlorate containing 4.5 gm. of potassium chloride. 

11. Iron and sulpur unite in the ratio of 7:4. Write the 
equation for the reaction. What weight of the product will be 
formed from (a) 20 gm. of iron and (b) 20 gm. of sulphur ? 



114 INORGANIC CHEMISTRY 

12. What weight of (a) zinc and (b) hydrochloric acid are needed 
to produce 17.5 gm. of hydrogen? 17.5 1. (at 0° C. and 760 mm.) ? 

13. How much ferric oxide will yield 2.5 gm. of iron? 

14. How much silver chloride can be produced from silver 
nitrate and 22.8 gm. of crystallized barium chloride (BaCl 2 . 2 H 2 0) ? 

15. Calculate the total weight of water that can be obtained 
from a metric ton of crystallized calcium sulphate. 

16. A balloon holds 150 kg. of hydrogen. How much (a) zinc 
and (b) sulphuric acid are needed to produce the gas? 

17. If water and 10 gm. of sodium interact, calculate the weight 
of each product. 

18. In the reduction of copper oxide by hydrogen 2.52 gm. of 
the solid product resulted. What weights of copper oxide and 
hydrogen were used ? 

19. A liter of oxygen (at 0° C. and 760 mm.) was transformed 
by phosphorus into phosphorus pentoxide. How many grams 
of phosphorus and of phosphorus pentoxide were involved ? 

20. A lump of carbon weighing 10 gm. is burned in air. What 

(a) volume and (b) weight of carbon dioxide is formed ? 

21. What weight of iron oxide (Fe 3 4 ) is formed by burning 
a metric ton of iron in oxygen? What volume of oxygen (at 0° C. 
and 760 mm.) is used? 

22. A lump of sulphur weighing 12 gm. is burned in air. Calcu- 
late (a) the weight and (b) the volume of oxygen needed and sul- 
phur dioxide formed. If air contains 21 per cent of oxygen (by 
volume), what volume of air is used? 

23. Calculate the weight of oxygen needed to burn 33 gm. of 
magnesium containing 11 per cent of impurities. 

24. A liter of hydrogen (at 0° C. and 760 mm.) is produced 
by the interaction of aluminium and sodium hydroxide. What 
weight of the other substances are involved? 

25. Write the equation for the combination of sulphur and 
oxygen when they unite in the ratio (by weight) (a) 1 : 1 and (6)2:3; 
for carbon and oxygen in the ratio (a) 3 : 8 and (b) 3 : 4 ; for nitro- 
gen and hydrogen in the ratio 4.66 : 1 ; for nitrogen and oxygen 
in the ratios 7 : 4 ; 3.5:8; 3.5:6; 3.5:10. 

26. Write equations for the following reactions : (a) Magnesium 
and sulphuric acid form magnesium sulphate and hydrogen. 

(b) Zinc sulphate and barium nitrate form barium sulphate and 
zinc nitrate, (c) Strontium carbonate and hydrochloric acid form 
strontium chloride, water, and carbon dioxide. 



CHAPTER VIII 
The Atmosphere — Argon and Related Elements — Nitrogen 

The atmosphere is the great envelope of gas surrounding 
the earth. It extends into space to an estimated height 
of fifty to two hundred miles. We live at the bottom of 
this vast ocean of air, as it is often called. 

Aristotle (384-322 b.c.) regarded air as one of the four elementary 
principles whose combinations made up all substances in the universe. 
The other three were earth, fire, and water. He taught that air pos- 
sesses two fundamental properties — heat and dampness. The early 
chemists used the word air in the sense in which the word gas is now 
employed. Thus, we have already learned that hydrogen was first 
called inflammable air. 

The terms atmosphere and air are often used interchange- 
ably, though by air we usually mean a limited portion of 
the atmosphere. Many skillful chemists have studied the 
action of air on living things, its relation to combustion, the 
effect of its weight, its composition, and its varied properties. 
Their work has contributed many fundamental facts to 
science. 

General Properties of the Atmosphere. — Air has w r eight. 
We often use the expression " light as air." But a cubic 
foot of air weighs 1.28 oz. and a room 40 X 50 X 25 ft. con- 
tains about two tons of air. A liter of dry, normal air at 
0° C. and 760 mm. weighs 1.293 gm. The total weight of 
the atmosphere has been estimated to be five thousand 
millions of millions of tons. The enormous mass resting 
upon the earth exerts a pressure w T hich is about fifteen pounds 
on every square inch. The amount of pressure upon a 

115 



116 INORGANIC CHEMISTRY 

square inch is called " an atmosphere/' and it is sometimes 
used as a unit of pressure; e.g. three atmospheres means 
a pressure of forty-five pounds per square inch. It is atmos- 
pheric pressure which causes water to rise in pumps and 
flow through siphons. Atmospheric pressure is exerted in all 
directions and is variable. It is measured by the barometer. 
The normal or standard pressure of the atmosphere, as al- 
ready stated, is equal to the pressure of a column of mercury 
which is 760 mm. (or 29.92 in.) high. In very accurate ex- 
periments certain mathematical corrections must be made 
in the height of the column as read on the barometer scale, 
but in ordinary work it is necessary to know the height only 
of the mercury column in order to know the pressure. The 
pressure of the atmosphere varies as the height and com- 
position of the atmosphere vary, and the barometer changes 
accordingly. 

Ingredients of the Atmosphere. — The atmosphere is a 
mixture of several gases. But since this mixture always 
contains approximately 78 parts of nitrogen and 21 parts 
of oxygen by volume, we often speak of air as consisting 
solely of these two gases. Besides this large proportion 
of oxygen and nitrogen, the air always contains small and 
variable proportions of water vapor and carbon dioxide gas. 
In addition to these four ingredients, air always contains 
the gas argon (and the related inert gases), and usually 
very small proportions of ozone, hydrogen, hydrogen di- 
oxide, compounds related to ammonia and nitric acid, dust, 
and germs. The composition varies but slightly in dif- 
ferent localities. Near the city air may contain a relatively 
larger proportion of dust, ammonia, sulphur compounds, 
and acids ; in the country the proportion of ozone is rela- 
tively large ; over the ocean and near the seacoast the air 
contains salt. 



THE ATMOSPHERE 1 17 

General Properties of Nitrogen. — The chemical element; 
nitrogen, which constitutes about 78 per cent of the at- 
mosphere (by volume), is a colorless gas, and has no taste 
or odor. It is somewhat lighter than air, and is slightly 
soluble in water. In many respects it differs markedly 
from oxygen. Thus, it will not support combustion, neither 
will it burn, nor sustain life. Animals die if left in nitrogen. 
Nitrogen is not poisonous, for the air we breathe contains 
a large proportion of nitrogen. Its function in the at- 
mosphere is to dilute the oxygen. It is an inert element 
compared with many others, although it combines directly 
with oxygen and a few other elements. At the ordinary 
temperature it is chemically indifferent, but at high tempera- 
tures it is quite active (see pages 129, 212, 217). 

The fact that a candle flame quickly goes out and a mouse soon 
dies in nitrogen was first observed by Rutherford, a Scottish physician, 
who discovered the gas in 1772. Soon after, Lavoisier showed the 
true relation of nitrogen to the atmosphere. To emphasize the ina- 
bility of the gas to support life, he called the new gas azote, the name 
now used for it by French chemists. 

Oxygen and Nitrogen in the Atmosphere. — The chemical 
activity of the atmosphere is due to the free oxygen it con- 
tains, as we have already learned in studying oxygen. If 
the air were largely oxygen, rusting and decay would proceed 
with astonishing rapidity, and fires once started would burn 
with great violence. On the other hand, nitrogen is chemi- 
cally inactive, and if the air contained much more than the 
normal amount, the chemical action of oxygen would be 
slower. Oxygen alone is too active, while nitrogen alone 
is rather inactive. To be serviceable to man, oxygen must 
be diluted with nitrogen, while nitrogen must be accom- 
panied by a small proportion of oxygen. 

The presence of oxygen and nitrogen in the atmosphere and the 
functions of the two gases were first clearly explained by Lavoisier 



118 



INORGANIC CHEMISTRY 



in 1777, though many others — Boyle, Priestley, Rutherford, and 
Scheele — helped solve the problem. 

Composition of the Atmosphere. — Samples of air from 
various parts of the globe show a remarkable uniformity 
of composition. For many years it was believed that pure 
air consisted solely of oxygen and nitrogen. But in 1895 it 
was found that nearly 2 per cent (by weight) of the gas 
hitherto called nitrogen is argon. (See Argon, below.) Ac- 
cording to the most recent results, the following is — 

The Composition of Pure Normal Dry Air 



Ingredient 


Percentage 




By Volume 


By Weight 


Nitrogen 

Oxygen 


78.122 

20.941 

0.937 


75.539 
23.024 


Argon . 


1.437 



The composition of the atmosphere was studied by Priestley, 
but his results were conflicting. Cavendish, in 1781, was the first to 
show that the proportion of oxygen and nitrogen in air is nearly 
constant. Since his time this result has been confirmed by many 
chemists, especially by Bunsen. In recent years the composition of 
the atmosphere and the properties of its inert components have been 
assiduously studied by Ramsay and others. 

The Volumetric Composition of Air can be found in the 
laboratory by introducing a known volume of pure air into 
a eudiometer and exploding it with a known and sufficient 
volume of hydrogen. The nitrogen does not participate 
in the chemical change, but all the oxygen in the air combines 
with twice its volume of hydrogen, forming a minute quantity 
of water; hence one third of the diminution in volume is 
the volume of oxygen in the air. The difference between 



THE ATMOSPHERE 



119 



the volume of oxygen found and the original volume of air 
is the volume of the other constituents — chiefly nitrogen. 
An illustration will make this experiment clear. Suppose 
(1) we mix and explode 100 cc. of air and 50 cc. of hydrogen, 
or 150 cc. in all, and (2) the residue measures 87 cc. Now, 
150 — 87 = 63, hence 63 cc. of the total volume of the mix- 
ture combined to form water (the 63 consisting of all the oxy- 
gen and part of the hydrogen). But one third of the 63 cc. is 
the oxygen which was in the original volume of air, because 
oxygen and hydrogen unite volumetrically in the ratio of 1 
to 2. Hence, 63 -f- 3 = 21, the volume of oxygen in 100 cc. 
of air. The remainder, 79 cc, is nitro- 
gen, argon, and other gases. 

Another Method, often used to determine 
the volumetric composition of the air, is based 
on the fact that phosphorus combines slowly 
with oxygen, even at the ordinary tempera- 
ture. The operation is performed by insert- 
ing a piece of phosphorus (see Figure 13 a) 
into a graduated glass tube containing a meas- 
ured volume of air. White fumes indicate 
immediate action. These fumes are solid par- 
ticles of phosphorus pentoxide. They soon 
dissolve in the water, which rises higher in 
the tube, as the oxygen combines with the 
phosphorus. In a few hours the phosphorus 
is removed, and the volume of gas is read. 
The difference between the first and last vol- 
umes is oxygen. The gas remaining in the 
tube is a mixture of nitrogen and argon. In 
performing this experiment unusual care 
must be taken not to touch the phosphorus 
with dry hands. Both gas volumes should 
be corrected for pressure, temperature, and aqueous vapor. 

The Gravimetric Composition of Air, as already stated, 
is : — 




Fig. 13 a. — Apparatus 
for determining the 
volumetric composi- 
tion of air. 



120 INORGANIC CHEMISTRY 

Oxygen 23.024 parts by weight, 

Nitrogen . . . . . . 75.539 parts by weight, 

Argon 1.437 parts by weight. 

Dumas and Boussingault in 1841 made the first accurate deter- 
mination of the gravimetric composition of the air. They passed 
pure air through a weighed tube containing copper, and arranged 
so that heat could be applied. The oxygen of the air combined 
with the copper, while the nitrogen passed on into a weighed globe. 
Both tube and globe increased in weight. The increase in the tube 
was the weight of the oxygen, while the increase in the globe was 
the weight of the nitrogen. 

Water Vapor in the Atmosphere. — Water vapor is always 
present in the atmosphere, owing to constant evaporation 
from the ocean, other bodies of water, and the soil. The 
total amount is large, though variable. A given volume of 
air at a given temperature will absorb a definite volume of 
water vapor and no more. The amount absorbed depends 
largely upon the temperature. Air containing its maximum 
amount of water vapor at a given temperature is said to be 
saturated at that temperature, or to contain 100 per cent 
of water vapor. The saturation point is also called the 
dew-point. On a pleasant day, the relative humidity, i.e. 
the relative amount of water vapor present, may vary from 
30 to 90 per cent, the average being about 50 per cent, 
though it varies with the locality. Warm air holds more 
vapor than cool air. The amount of water vapor in the air 
has a marked influence on. the physical condition of man. 
The depressing weather during " dog days " is due to the 
high relative humidity of the air, which sometimes reaches 
nearly 100 per cent. The specialized forms of life, both 
animal and vegetable, found in deserts are largely due to 
the dry air. The languor felt in a " close " room or crowded 
hall is partly caused by the excess of water vapor in the 
" bad " air. The presence of water vapor in the air is shown 



THE ATMOSPHERE 121 

by the moisture which collects on the outside of a vessel 
containing cold water, such as a pitcher of iced water. The 
moisture comes from the air around the vessel. For a 
similar reason, water pipes in a cellar and the cellar walls 
themselves are moist in summer. The deliquescence of 
calcium chloride, common salt, and other substances like- 
wise reveals the presence of water vapor in the air. (See 
Deliquescence.) When the temperature of the air falls 
sufficiently, the water vapor condenses and is deposited in 
the form of dew, rain, fog, mist, frost, snow, sleet, or hail. 
The clouds are masses of water vapor which has been con- 
densed by the cold of the upper air. 

Carbon Dioxide in the Atmosphere. — Carbon dioxide 
is one product of the respiration of animals and of the 
combustion and decay of organic substances. By these 
processes an immense quantity of carbon dioxide is being 
constantly poured into the atmosphere. The proportion in 
the atmosphere is variable, though not between such wide 
limits as the water vapor. The proportion in normal air 
is 3 to 4 parts in 10,000 parts of air. Over the ocean the 
proportion is smaller, but in the air of cities it is greater. 
In crowded rooms the proportion is often as high as 33 
parts in 10,000. The proportion of carbon dioxide in the 
atmosphere as a whole is practically constant, largely owing 
to the fact that this gas is an essential food of plants. (See 
Carbon Dioxide.) The presence of carbon dioxide in the 
air is detected by calcium hydroxide. If a solution of 
calcium hydroxide is exposed to the air, the carbon dioxide 
interacts with the calcium hydroxide, forming a thin, white 
crust of insoluble calcium carbonate on the surface of the 
liquid. If air is drawn through calcium hydroxide, the liquid 
becomes milky, because the fine particles of calcium car- 
bonate remain temporarily suspended in the liquid. The 



122 INORGANIC CHEMISTRY 

purity of air is often determined by finding out what pro- 
portion of carbon dioxide it contains. If a known volume 
of dry air is drawn through a known weight of calcium 
hydroxide or similar liquid, the increase in weight will be 
the weight of carbon dioxide in the volume of air used. 
The equation for the interaction of carbon dioxide and 
calcium hydroxide is : — 

C0 2 + Ca(OH) 2 = CaC0 3 + H 2 

Carbon Calcium Calcium Water 

Dioxide Hydroxide Carbonate 

The different gases in the atmosphere are not arranged 
in layers according to their densities. They are in constant 
circulation. (See Diffusion.) Hence carbon dioxide, though 
heavier than oxygen and nitrogen (volume for volume), 
does not remain nearest the ground, but is distributed 
through the air. In a few exceptional localities, carbon 
dioxide arises from volcanic openings faster than it can 
diffuse, and fills the cave or adjacent valley. 

Argon (A) and Related Elements. — Argon, as already 
stated, is an essential and constant component of the atmos- 
phere. Argon was discovered in 1895 by Rayleigh and Ram- 
say. Rayleigh found that nitrogen extracted from air had 
a greater weight than an equal volume of nitrogen obtained 
from compounds of nitrogen. Consequently, they believed 
that the nitrogen from air contained another gas hitherto 
overlooked. A series of elaborate experiments showed 
that after the oxygen and nitrogen were removed from 
purified air, there still remained a small quantity of a new 
gas, which they called argon. It may be obtained (1) by 
passing pure air over heated copper to remove the oxygen, 
and then the remaining gas over heated magnesium or 
calcium to remove the nitrogen; or (2) by passing electric 
sparks through a mixture of air and oxygen, and removing 




SIR WILLIAM RAMSAY 



THE ATMOSPHERE 123 

the compound of oxygen and nitrogen as fast as it is formed. 
The latter method is a repetition of the one used by Cavendish 
when he determined the composition of air, and he would 
have no doubt discovered argon had he continued his 
investigations. As stated above, the proportion in the 
atmosphere is .937 per cent by volume and 1.437 per cent 
by weight. 

Argon is a colorless, odorless gas. It dissolves in water to 
the extent of about 4 volumes in 100. One liter of argon at 
0° C. and 760 mm. weighs 1.7809 gm. (compare with nitrogen, 
page 129). Liquid argon boils at — 186° C. and solid argon 
melts at —189.5° C. A conspicuous property of argon is 
its lack of chemical activity. No compounds of this element 
have as yet been prepared or discovered. The name argon 
is happily chosen, being derived from Greek words signify- 
ing inert. 

Helium (He), neon (Ne), krypton (Kr), and xenon (Xe) 
are inert gases discovered by Ramsay subsequently to argon. 
They constitute an exceedingly minute proportion of the 
atmosphere. Like argon, they do not form compounds. 
Their proportions in the atmosphere are approximately : — 

Helium, 3 to 4 parts per million, 
Neon, 1 to 2 parts per hundred thousand, 
Krypton, 1 part in 20,000,000, 
Xenon, 1 part in 170,000,000. 

Helium was detected in the atmosphere of the sun by 
Lockyer in 1868. It was found by Ramsay, soon after he 
discovered argon, in the gases expelled from certain minerals 
and in the gas and water of certain mineral springs. Helium 
is one of the disintegration products of radium. Niton (Nt) 
belongs to this group of elements (see Radioactivity). 

The process of separating the inert gases of the atmosphere 
from each other is an excellent illustration of modern ex- 



124 INORGANIC CHEMISTRY 

perimental methods. The mixture of the five gases is. 
compressed in a bulb and cooled to about — 185° C. by 
immersion in liquid air; the argon, krypton, and xenon 
condense to a liquid in which the helium and neon dissolve. 
When the bulb is removed and warmed, the helium and neon 
together with considerable argon escape first into a special 
bulb, the argon next, and finally the krypton and xenon. 
Several repetitions, however, are necessary to separate the 
argon from the helium and neon and the krypton and xenon, 
as well as the last two from one another. By immersing the 
bulb containing the helium and neon in liquid hydrogen, 
the neon solidifies and the helium can be removed first by 
a pump; subsequently the neon when warmed can be simi- 
larly removed as the pure gas. 

When these inert gases are examined by a spectroscope, 
they exhibit striking spectra. That is, when electric sparks 
are passed through a closed tube containing any of these 
gases and the light thereby produced is viewed through a 
spectroscope, many colored vertical lines are seen. Certain 
lines are conspicuous, e.g. the orange line in the case of 
helium, and by means of these and other lines it is possible 
to detect small quantities of these gases and to distinguish 
them from other gases, since no two spectra are exactly 
alike under given conditions. ' (For an account of the spec- 
troscope and its application, see Chapter XXV.) 

Air is a mixture, in spite of the fact that we speak of its 
" composition." Chemical compounds, as we have already 
learned, have two invariable characteristics; viz. (1) their 
constituents are in a fixed proportion, and (2) their formation 
and decomposition are usually attended by definite evidences 
of chemical action, such as light, heat, electrical phenomena, 
change of color, etc. The following facts show that air is 
a mixture of free gases : — 



THE ATMOSPHERE 125 

(1) The proportion of oxygen and nitrogen in the air is 
not fixed, but varies, though between very narrow and deter- 
minable limits. 

(2) When nitrogen and oxygen are mixed in the propor- 
tions w 7 hich form air, the product is exactly like air, but the 
act of mixing gives no evidence of chemical action. 

(3) When air is dissolved in water, a greater proportion 
of oxygen than nitrogen dissolves; i.e. they dissolve as 
independent gases in proportions fixed by their intrinsic solu- 
bility and partial pressure. (See Solubility of Gases.) If the 
oxygen and nitrogen were combined in the air, the dissolved 
air would, of course, have the same composition as air 
itself. 

(4) Oxygen and nitrogen distill independently from 
liquid air. 

Liquid air is a mixture of the liquefied gases which con- 
stituted the air used. It is a milky liquid, owing to the 
presence of solid carbon dioxide and ice. If these solids 
are removed by filtering, the filtrate has a pale blue tint. 
It is intensely cold, and boils at about — 190° C. under 
atmospheric pressure. If a vessel is filled with liquid air, 
the latter boils vigorously, the surrounding air becomes 
very cold, frost gathers on the vessel, and in a short time 
the liquid air will have entirely disappeared into the air of 
the room. If, however, the liquid air is placed in a Dewar 
bulb, the evaporation is only slightly affected by changes 
of temperature. 

The Dewar bulb (Fig. 14) consists of two flasks, one 
within the other, sealed together by an air-tight joint at the 
top; the space between the flasks is a vacuum. Sometimes 
the surfaces of the flasks are coated with silver, which re- 
flects the heat and thereby retards the evaporation of 
the contents, 



126 



INORGANIC CHEMISTRY 



Liquid air, owing to its extremely low temperature, 
produces remarkable physical changes. A tin or iron vessel 
which has been cooled by liquid air is so brittle that it may 
often be crushed with the fingers. Nearly all plastic or 
soft substances, including many kinds of food, when im- 
mersed in liquid air become hard and brittle, leather being 




k^ 




Fig. 14*. — Dewar bulbs. 

the only important exception. Mercury freezes so hard 
in liquid air that it can be used as a hammer to drive a nail. 
When liquid air is put in a teakettle standing on a block 
of ice the liquid air boils vigorously. If the kettle of liquid 
air is placed over a lighted Bunsen burner, frost and ice 
collect on the bottom of the kettle, because the intense cold 
of the kettle solidifies the water vapor and carbon dioxide 
which are the two main products of burning illuminating 
gas. If water is now poured into the kettle, the liquid 
air boils vigorously and the water is quickly frozen; the 
water is so much hotter than' the liquid air that the latter 
boils more violently, and since its rapid evaporation causes 
absorption of heat, the water loses its heat and becomes 
ice. Ordinary liquid air is from one fifth to one half liquid 
oxygen, and will support combustion. A red-hot rod of 
steel or of carbon burns brilliantly in this cold liquid. 



THE ATMOSPHERE 127 

Numerous applications of liquid air have been proposed, 
but thus far they have not passed the experimental stage. 
It has been suggested that it be used as a refrigerant instead 
of ice, for ventilating and cooling rooms, as a blasting 
material, for removing diseased flesh from a wound, for de- 
stroying refuse, and as a source of oxygen and nitrogen. The 
last use is based primarily on the fact that as liquid air 
evaporates the nitrogen passes off first, and in a short time 
relatively pure oxygen remains. (See Oxygen.) 

A little liquid air was produced in 1883 w r ith considerable 
labor and at an enormous expense. At present it is easily 
manufactured in large quantities at a comparatively low 
cost. Compressed air cooled by water is forced through 
a pipe to a valve. As it escapes through the valve, it expands 
and its temperature falls, because expansion is a cooling 
process. After expansion, the cold air is led back over the 
outer surface of the same pipe by which it came, whereupon 
it rapidly regains its former temperature. But in doing 
so it cools the pipe itself and the air within it. This latter 
air in turn expands and falls in temperature, but as it was 
cooler than the first portion before expansion, so it is colder 
than it after expansion. As the pressure within the pipe 
is maintained by a continuous supply of compressed air, 
the pipe becomes continually colder until finally the expand- 
ing air at the valve liquefies in part and is collected in a 
suitable receptacle. 

Liquefaction of Gases. — In the first methods used to 
liquefy gases, the gas was subjected simultaneously to a 
great pressure and a low temperature. Thus, Faraday 
about 1823 liquefied chlorine gas by heating one arm of a 
sealed bent tube containing a chlorine compound after 
having placed the other end in a freezing mixture ; the 
liberated chlorine being unable to escape was liquefied by 



128 INORGANIC CHEMISTRY 

the pressure and low temperature. Other gases were 
liquefied by a similar method. A few, however, could not 
be liquefied, e.g. oxygen, hydrogen, and nitrogen, and these 
were called permanent gases. About 1870 it was shown 
that if these so-called permanent gases were cooled to a 
sufficiently low temperature, they could be liquefied if the 
pressure was also sufficiently great. It was found, further- 
more, that each gas has a critical temperature, i.e. a tem- 
perature below which it must be cooled to produce lique- 
faction, no matter what the pressure. Thus, the critical 
temperature of oxygen is about —118° C, that of carbon 
dioxide is about + 31° C, that of normal air is about —.140° 
C, and that of sulphur dioxide is about + 155.5° C. Ob- 
viously, critical temperatures vary between wide limits. 
The pressure which must be applied to liquefy a gas at its 
critical temperature is called its critical pressure. The 
latter value varies, but not between such wide limits as the 
critical temperature. It is about 58 atmospheres for oxygen, 
15.3 for hydrogen, and 113 for ammonia (NH 3 ). As the 
temperature falls below the critical point, less pressure than 
the critical amount is needed for liquefaction, and if the 
temperature of the gas is reduced to the boiling point of the 
liquefied gas, no external pressure whatever is needed for 
liquefaction. Hence the essential point in liquefying most 
gases is the production of a sufficiently low temperature. 
This can be done in some cases by external application of 
cold, though in the case of gases having a low critical tem- 
perature/the cooling is now produced by a purely mechanical 
process, e.g. like that used for liquefying air (see above). By 
this process all known gases have been liquefied. 

If a liquefied gas can be cooled to a sufficiently low tem- 
perature, it becomes solid. Thus, Dewar by boiling liquid 
hydrogen under reduced pressure froze it to a foam-like 
solid, the temperature being about —258° C. 



THE ATMOSPHERE 129 

Nitrogen 
Occurrence. — Nitrogen, besides comprising four fifths 
of the atmosphere, is a constituent of nitric acid and am- 
monia, and of many compounds related to them. It is 
also an essential constituent of animal and vegetable matter. 

The name nitrogen was given to the gas by Chaptal from the fact 
that it is a constituent of niter, an old name of potassium nitrate. 

Preparation. — Nitrogen can be obtained from the air 
by removing the oxygen by phosphorus. A tall jar is 
placed over burning phosphorus contained in a shallow 
dish floating in a large vessel of water. The oxygen com- 
bines with the phosphorus, leaving nitrogen, more or less 
pure, in the jar. Other methods may be used, such as 
passing air over heated copper, or decomposing ammonium 
dichromate by heat. It is prepared in the laboratory by 
heating a mixture of sodium nitrite and ammonium chloride. 
It can be prepared commercially from liquid air (see page 17). 

Additional Properties. — In addition to its inertness, 
already mentioned, nitrogen is a little lighter than air, 
and is very sparingly soluble in water. Its density is .972 
(air = 1). One liter at 0° C. and 760 mm. weighs 1.2507 
gm. One hundred liters of water dissolve only about 1.5 1. 
at the ordinary temperature. The critical temperature 
is about —146° C. Liquid nitrogen boils at —195.5° C. 
under ordinary atmospheric pressure, and solid nitrogen 
melts at about -214° C. 

It combines directly with silicon and also with many metals 
at a red heat, forming nitrides, e.g. magnesium nitride 
(Mg 3 N 2 ) . At high temperatures and under special conditions 
nitrogen combines with oxygen and with hydrogen, forming 
nitric oxide (NO) and ammonia (NH 3 ). (See pages 212, 
217.) 



130 



INORGANIC CHEMISTRY 



Relation of Nitrogen to Life. — Oxygen, carbon dioxide, 
and water vapor are essentially related to the life of plants 
and animals. Nitrogen is also vitally connected with 
different forms of life. Atmospheric nitrogen merely dilutes 
the oxygen. Although we live in an atmosphere contain- 
ing such a large proportion of nitrogen, we cannot assimilate 
it. According to a reliable authority, " the air as it leaves 
the lungs contains 79.5 per cent of nitrogen," and hence 
cannot become a part of the body. Yet all flesh contains 
nitrogen, and certain rejected waste products of animals 




A B 

Fig. 15. — A leguminous plant (the hairy vetch) with (B) and without 
(A) nodules on the roots. 



contain considerable combined nitrogen. The nitrogen 
needed by animals must be in combination to become avail- 
able. And it is taken in the form of nitrogenous food 
such as lean meat, fish, wheat and other grains. 

Most plants take up combined nitrogen from the soil in 
the form of nitrates (compounds derived from nitric acid) 
or of ammonia. Hence combined nitrogen is being con- 



THE ATMOSPHERE 131 

stantly taken from the soil, and in order to preserve the fer- 
tility of the soil, nitrogen must be supplied. This is done 
by adding to the soil a fertilizer containing nitrogen com- 
pounds. Sometimes nitrogenous organic matter is used, 
such as manure, dried blood, and meat or fish scraps. Chem- 
ical fertilizers are extensively used, e.g. sodium nitrate 
(NaN0 3 ), ammonium sulphate ((NH 4 )2S0 4 ), calcium nitrate 
(Ca(N0 3 )2), or calcium cyanamide (CaN 2 C). Leguminous 
plants, such as peas, beans, and clover, assimilate free nitro- 
gen directly from the air by means of bacteria which are in 
nodules on their roots (Fig. 15). This process is called fixa- 
tion of nitrogen. Sometimes soils are treated with a prep- 
aration which contains " nitrogen bacteria. " 

Problems and Exercises 

1. A quantity of air measures 24 1. at 15° C. ; the temperature 
is reduced to — 16° C. What is now the volume? (Pressure un- 
changed.) 

2. If 100,000 cu. m. of air at any pressure were saturated with 
moisture at 20° C, what weight of water is deposited at 0° C? 

3. If air, at 760 mm., has a temperature of 20° C, and its dew- 
point, i.e. the temperature at which it is saturated with moisture, 
is 15° C, what per cent of moisture by volume does it contain? 

4. If 1 cc. of dry air, under standard conditions, weighs .00129 
gm., what would be the weight of 1 cc. of air saturated with mois- 
ture at 30° C. under normal pressure? 

5. How many (a) grams and (6) cubic centimeters of argon can 
be obtained from 1500 kg. of pure dry air? (Standard con- 
ditions.) 

6. How many grams of copper oxide (CuO) are formed by pass- 
ing 1728 gm. of normal air over pure copper? How many grams 
of magnesium nitride (Mg 3 N 2 ) by passing the residual gas over red- 
hot magnesium? How many grams of gas remain? Equations 
are Cu + O = CuO and 3 Mg + 2 N = Mg 3 N 2 . 

7. Write equations for the following reactions : (a) Phosphorus 
pentoxide and water form phosphoric acid (H 3 P0 4 ). (6) Copper 
oxide (CuO) and hydrogen from copper and water. 



CHAPTER IX 
Solution — Theory of Electrolytic Dissociation 

Introduction. — Many facts about solutions have already 
been stated (see Chapter V). The present chapter is a con- 
tinuation of the discussion of solutions with special reference 
to the theory of their nature and the interpretation of certain 
phenomena by this theory. The term solution will be re- 
stricted to aqueous solutions, i.e. those in which water is the 
solvent. 

General Properties of Solutions. — It is very desirable to 
recall at this point certain properties of solutions. First, 
the solubility of the solute in the solvent varies widely and is 
generally increased by rise of temperature until it reaches a 
limit, thereby giving the phenomena connected with unsatu- 
ration, saturation, and supersaturation. Second, the solute 
in most cases can be recovered unchanged by evaporating, 
cooling, or distilling the solution. Third, the solution often 
contains the solute in an especially favorable condition for 
chemical action. Finally, solutions have a definite boiling 
point, freezing point, and vapor pressure, which differ from 
the corresponding values of the solvent. 

Behavior of Solutions toward an Electric Current. — Pure 

water and pure dry, solid sodium chloride do not appreciably 
permit the passage of an electric current. But a solution of 
sodium chloride is an excellent conductor of electricity, and 
the same is true of a solution of hydrochloric acid and of 
sodium hydroxide. On the other hand, sugar is a non-con- 
ductor, both in solution and in the solid state. In a word, 

132 



SOLUTION 133 

solutions of certain substances conduct electricity, while 
solutions of others do not. Hence solutions can be divided 
on this basis into two classes, viz. : (1) electrolytic solutions, 
or those which conduct electricity, and (2) non-electrolytic 
solutions, or those which do not conduct electricity. Sub- 
stances whose solutions are electrolytic are called electro- 
lytes, and those whose solutions are not electrolytic are called 
non-electrolytes. Each class has characteristic properties. 
But electrolytes possess certain conspicuous properties 
which are not exhibited by non-electrolytes, and it was the 
study of these characteristic properties that led to the pro- 
posal of the present theory of solutions. It is called the 
theory of electrolytic dissociation and was proposed by the 
Swedish physicist Arrhenius in 1887; its general adoption 
has been hastened by the work of Van't Hoff, Ostwald, and 
Nernst. 

Theory of Electrolytic Dissociation. — It was believed for 
many years that the molecules of a dissolved substance w^ere 
distributed unchanged throughout the solvent. It was 
also believed that the molecules of certain dissolved sub- 
stances combined to some degree with the molecules of the 
solvent. Evidence is fast accumulating which indicates that 
in many solutions the solute is in the form of molecules, 
while in others the solvent does unite with the solute or some 
of its constituents. The present theory of solutions differs 
from these by offering an explanation of solution which is 
more comprehensive. Briefly, the theory of electrolytic dis- 
sociation assumes (1) that molecules of electrolytes when 
dissolved in water break up to a varying degree into inde- 
pendent particles charged with electricity, and (2) that the 
nature and number of these electrically charged particles 
determine to a large extent certain physical and chemical 
properties of solutions. 



134 INORGANIC CHEMISTRY 

Before stating the facts on which the theory is based, it 
will be necessary to expand the two assumptions and to de- 
fine several terms. The breaking up of certain substances 
when in aqueous solution is called electrolytic dissociation, 
or ionization. The independent particles are called ions. 
Thus, when sodium chloride is dissolved in water, some of its 
molecules dissociate into sodium ions and chlorine ions. 
Each ion is a portion of a molecule and is charged with elec- 
tricity. Two kinds of ions are present in every electrolytic 
solution, viz. electro-positive ions, or cations, and electro- 
negative ions, or anions. Ions, although formed by the 
dissociation of molecules, are not identical with atoms, but 
differ mainly in having a charge of electricity. For exam- 
ple, when sodium chloride is dissolved in water, the electro- 
positive sodium ions move about in the water without pro- 
ducing any apparent chemical change; but ordinary sodium 
interacts violently with water, as we have already seen. 
Similarly, the chlorine ions circulate freely in water and ex- 
hibit none of the effects of gaseous chlorine upon water. In 
a word, in such a solution the sodium ions and chlorine ions 
exist side by side without any apparent decomposition of the 
water or any apparent tendency to combine with each other. 
The properties of ions, as already stated, are mainly due to 
their electric charges, and ions may be defined as electrically 
charged atoms or atomic groups. It is customary to repre- 
sent ions by chemical symbols supplemented by the sign 
which designates the kind of electric charge. Thus, the 
ions formed by the dissociation of sodium chloride are Na" 1 " 
and CI", while potassium nitrate yields K + and N0 3 ". A 
solution of sodium chloride gives no evidence of electricity. 
In general, solutions of electrolytes are electrically neutral; 
i.e. although they allow an electric current to pass through 
them when supplied from some external source such as a 
battery or dynamo, the solution itself is electrically neutral. 



SOLUTION 135 

It therefore follows that the sum of the electric charges of 
the positive ions equals the sum of the electric charges of 
the negative ions. The ions balance each other electrically. 
Thus, in the case of sodium chloride solution, the number 
of sodium ions equals the number of chlorine ions and the 
sum of the positive charges on the sodium ions equals 
the sum of the negative charges on the chlorine ions. In the 
case of calcium chloride (CaCl 2 ), each molecule dissociates 
into two chlorine ions and one calcium ion; but since the 
sum of each kind of electric charges must be the same, each 
calcium ion must carry twdce the charge which is on each chlo- 
rine ion. Hence the calcium ions are designated as Ca ++ and 
the chlorine ions as 2 CI". A unit charge of electricity is 
indicated by the single sign, and multiples by the proper 
number. For example, the sodium ion is Na + , the common 
copper ion is Cu ++ , the aluminium ion is Al +++ , the 
sulphate ion is S0 4 , and the nitrate ion is N0 3 ^. The 
degree of dissociation varies with the concentration and with 
the electrolyte. The greater the dilution, the greater the 
dissociation. Conversely, dissociation is decreased by de- 
creasing the volume of the solvent; the ions tend to unite, 
forming undissociated molecules which can be ultimately 
obtained as a mass of the original substance by evaporation. 
Consequently, a solution contains undissociated molecules 
as w T ell as ions, depending upon the substance and the con- 
centration of the solution. Experiment shows that the chem- 
ical behavior of a dissolved substance often depends largely 
on the extent of the dissociation of the molecules into ions. 

As previously stated, only certain substances are electro- 
lytes. These are acids, bases, and salts. The general prop- 
erties of these substances are discussed in Chapter X and 
their special characteristics are treated under the individual 
compounds. It is sufficient for our present purpose to re- 
gard them as a single class of substances which dissociate 



136 INORGANIC CHEMISTRY 

in solution into ions. That is, ions are in the solutions of the 
familiar acids like sulphuric, hydrochloric, and nitric, the 
familiar bases like sodium hydroxide and potassium hydrox- 
ide, and the familiar salts like potassium chlorate, sodium 
chloride, sodium sulphate, and many others which will soon 
be described. Sugar, alcohol, and other compounds not so 
familiar do not dissociate into ions when dissolved in water 
and their solutions do not conduct electricity. Neverthe- 
less, such solutions have specific and instructive properties, 
especially when compared with the properties of electro- 
lytic solutions. 

Properties of Electrolytes and Non-electrolytes. — The 

theory of electrolytic dissociation, which has just been dis- 
cussed, is based on many facts which have accumulated as 
the outcome of a comprehensive study of the properties of 
solutions, both electrolytic and non-electrolytic. These facts 
will now be presented. Careful distinction should be drawn 
between the statements in the paragraphs immediately pre- 
ceding and those about to be made. The preceding con- 
cerned the theory of electrolytic dissociation. The forth- 
coming concern facts and laws derived from experiment and 
interpreted by the theory. 

(a) Osmotic Pressure. — The passage of a liquid through a 
membrane is called osmosis. Certain membranes permit 
the passage of the solvent but prevent more or less the pas- 
sage of the solute ; such membranes are said to be semiper- 
meable. Osmosis is a common phenomenon in physiological 
processes, for semipermeable membranes occur in animal and 
vegetable organisms. Osmosis and osmotic pressure can be 
demonstrated by a simple experiment. A piece of parchment 
paper is tied tightly over the larger end of a thistle tube, 
which is then partly filled with a concentrated sugar solution 
and immersed in a vessel of water (Fig. 15a). The membrane 



SOLUTION 



137 




is permeable to water, but not to sugar. Water passes 
through the membrane into the sugar solution, which in- 
creases in volume and hence rises in the tube. If the mem- 
brane is strong enough, the column of solution will rise to a 
maximum height. At this point the 
weight of the solution, and therefore 
its pressure, is such that the tendency 
of the water to pass through the mem- 
brane into the solution equals the ten- 
dency of the water to pass out. When 
this condition is reached, the weight of 
the liquid above the membrane is a 
measure of the osmotic pressure of the 
sugar solution. 

In accurate measurements of the 
osmotic pressure of aqueous solutions the 
semipermeable membrane is usually a 
film of cupric ferrocyanide (Cu 2 Fe(CN) 6 ) FlG - 15 a - — Experiment 

j .. i • ,i r i i to illustrate osmotic 

deposited in the pores ol an unglazed 

^ ^ ° pressure. 

porcelain vessel. The cell, as it is called, 
is filled with the solution to be investigated, fitted tightly 
with a special form of stopper, connected with a manometer 
(to measure the pressure), and then immersed in water. 
Water flows slowly through the membrane into the cell 
until a maximum pressure is produced ; the system is then in 
equilibrium, i.e. the tendency of the water to pass into the 
solution through the semipermeable membrane is balanced 
by the opposing pressure of the manometer. This increase 
over the original hydrostatic pressure is read on the manom- 
eter and is equal to the osmotic pressure of the sugar solu- 
tion. Determinations of the osmotic pressure of dilute solu- 
tions of non-electrolytes show that osmotic pressure is 
directly proportional (1) to the concentration and (2) to the 
absolute temperature of the solution; in (1) the tempera- 



138 INORGANIC CHEMISTRY 

ture must be kept constant and in (2) the volume must be 
kept constant. Furthermore, if a gram-molecular weight 
(that is, the number of grams numerically equal to the mo- 
lecular weight) of non-electrolytes is dissolved in equal quan- 
tities of water, the osmotic pressure of each solution is the 
same. The gram-molecular weight of a substance is called a 
mole — a convenient value which is frequently used in stating 
concentration and in comparing experimental results. In the 
case of osmotic pressure, for example, one mole (342 gm.) of 
cane sugar (C12H22O11) and one mole (58 gm.) of acetone 
(C 3 H 6 0) dissolved in equal quantities of water have the same 
osmotic pressure. 

The osmotic pressure of an electrolytic solution is found 
to be greater than that of a non-electrolytic solution 
under the same conditions. The excess of pressure varies 
somewhat with the conditions, i.e. with concentration, tem- 
perature, and the substance. Nevertheless, the difference 
between the normal and the abnormal osmotic pressure is 
conspicuously large in the case of many substances yielding 
electrolytic solutions, these substances being acids, bases, 
and salts. Thus, a solution of potassium chloride has an 
osmotic pressure 1.88 times that of a corresponding sugar 
solution, while a calcium chloride solution has a pressure 
about 2.5 times. This discrepancy between the values of the 
osmotic pressure of non-electrolytic and electrolytic solutions 
is readily explained, if the facts are interpreted by the theory 
of electrolytic dissociation. According to this theory a non- 
electrolytic solution contains only molecules and an elec- 
trolytic solution contains both ions and molecules. Now 
osmotic pressure is believed to be due to the independent 
particles of the solute in solution, and the amount of pressure 
is determined by the number of particles in a given volume. 
Hence, when equivalent solutions are used, the electrolytic 
solution contains more particles than the non-electrolytic 



SOLUTION 



139 



because some of the molecules in the electrolytic solution 
are dissociated into ions. Facts and theory agree as far as 
osmotic pressure is concerned. 

(b) Freezing Point and Boiling Point of Solutions. — In 
the preliminary discussion of the properties of solutions given 
in Chapter V, it was stated that the freezing point of a solu- 
tion is lower than the freezing point of the solvent and that 
the depression of the freezing point is proportional to the 
concentration of the solution. These facts can be expressed 
by a concrete case, thus : — 

Depression of the Freezing Point 



Solvent 


Solute 


Freezing Point 


Depression 


100 gm. 

water 

in each case 


11.4 gm. sugar 
22.8 gm. 
34.2 gm. 


- .62° C. 
-1.23° C. 
-1.85° C. 


.62 

2 x .62 (approx.) 

3 x .62 (approx.) 



Furthermore, experiment shows that the freezing point of 
water is depressed the same number of degrees, if 1000 gm. of 
water contain a mole of certain solutes {i.e. a number of 
grams numerically equal to the molecular weight). Thus, 
the freezing point of water is depressed about 1.86° C. by a 
solution of 342 gm. of cane sugar (C^H^Ou) and 58 gm. of 
acetone (C 3 H 6 0), each dissolved in 1000 gm. of water. This 
uniform behavior is not well exhibited, however, unless the 
solutions are dilute and involve no chemical action between 
solvent and solute. On the other hand, when solutions of 
electrolytes, i.e. acids, bases, and salts, are experimented 
with, the freezing point is lower than that produced by non- 
electrolytes under the same experimental conditions. More- 
over, the depression is not uniform for all electrolytic solu- 
tions under uniform conditions. For example, a solution 
containing a mole of sodium chloride (i.e. 58.5 gm.) in 1000 



140 INORGANIC CHEMISTRY 

grams depresses the freezing point of water about 3.5° C, or 
nearly twice the amount produced by a cane sugar solution 
of equivalent concentration. This exceptional behavior of 
solutions of acids, bases, and salts can be explained as in the 
case of osmotic pressure. Solutions of non-electrolytes con- 
tain only molecules, while solutions of electrolytes contain 
ions into which some of the molecules have dissociated. 
Hence the number of independent particles (molecules and 
ions) in the electrolytic solution is greater than in the non- 
electrolytic solution. Ions and molecules act alike on the 
freezing point of a solution, and the larger the number of 
particles, the greater the depression. This deduction is 
further confirmed by the fact that electrolytic solutions in 
which a large proportion of molecules is dissociated show a 
relatively greater depression than those in which the disso- 
ciation is limited — the degree of dissociation being found by 
an independent experiment. 

Analogous statements can be made about the elevation of 
the boiling point of solutions. That is, (1) the boiling point 
of a solution is higher than that of the solvent, (2) the ele- 
vation of the boiling point is proportional to the concentra- 
tion, and (3) the elevation is the same (i.e. .52° C.) in the case 
of all non-electrolytic solutions containing a mole of the solute 
in 1000 gm. of water. But solutions of acids, bases, and salts 
behave exceptionally. They boil at a higher temperature 
than non-electrolytic solutions of the same concentration. 
The explanation offered by the theory of electrolytic disso- 
ciation is the same as in the preceding cases, viz. dissociation 
of some of the molecules into ions, which affect the boiling 
point in the same way as the molecules themselves. 

(c) Electrolysis of Solutions. — Electrolysis is the series 
of changes accompanying the passage of an electric current 
through a solution. It is accomplished in an electrolytic cell. 
This piece of apparatus consists of three essential parts — 



SOLUTION 



141 



the electrolytic solution, the containing vessel, and two 
electrodes which convey the electric current to and from the 
solution. A simple form of such a cell is shown in Figure 16. 
It is customary to speak of the 
current as entering the solution 
by the anode or positive electrode 
and as leaving by the cathode or 
negative electrode. Both the elec- 
trolytic cell and the ordinary vol- 
taic cell are fully described in 



g 




r* 


^^=^=^-^=^^^=^==^^1 


1 


n 




=■ 


===== =====- 


^^ 



Chapter XI, though a general idea Fig. 16. — Simple electrolytic 
of the electrolytic cell serves our celL A (anode) and C (cath " 



ode) are the electrodes which 
convey the current to and 
from the electrolytic solution 
(E) ; B or D is the battery 
or dynamo which provides 
the electricity. 



present need. 

When a concentrated solution of 
hydrochloric acid gas (commonly 
called simply hydrochloric acid) is 
put into an electrolytic cell and 

subjected to the action of an electric current, two gases 
are liberated, — hydrogen at the cathode and chlorine at 
the anode. This is a very simple illustration of electrolysis. 
Let us interpret it by the theory of electrolytic dissociation. 
When hydrochloric acid gas is dissolved in water, hydrogen 
ions (H + ) and chlorine ions (CI") are immediately formed 
by the dissociation of some of the molecules of hydrochloric 
acid (HC1). These ions are formed in the solution as soon 
as the acid dissolves and before the electric current is con- 
nected with the cell. As soon as the current is turned on, 
however, the electrodes at once become charged with elec- 
tricity — the anode assuming a positive charge and the 
cathode a negative charge. Now according to a principle 
established many years ago, bodies charged with like kinds 
of electricity repel each other and bodies charged with unlike 
kinds attract each other. Consequently the anions or electro- 
negative ions move toward the anode or electro-positive 



A is the anode and C is the cathode. 
B or D is the battery or dynamo. 



142 INORGANIC CHEMISTRY 

electrode, while the cations or electro-positive ions move 
toward the cathode or electro-negative electrode, or briefly, 
" anions to anode, cations to cathode. " This migration of 
the ions, as it is called, toward their respective poles is shown 
diagrammatically in Figure 17. As soon as the ions reach 

( their electrodes, they act in 

r^~ ' b or d ^"^ accordance with another long- 

established principle; that is, 
they give up their electric 
charges. In other words, when 
the electro-positive cations of 
hydrogen touch the electro- 

Fig. 17. — Migration of ions in an J . 

electrolytic cell. The cations are negative cathode, the electric 
marked H+ and the anions Cl~. charges are neutralized. Elec- 
tric charges are constantly re- 
newed on the cathode by the 
battery or dynamo, but the hydrogen ions once deprived of 
their electric charges do not regain them and immediately 
become ordinary, uncharged hydrogen atoms, which combine 
and escape as molecules of hydrogen gas. Similarly, the elec- 
tro-negative anions of chlorine migrate to the electro-positive 
anode, lose their charges, become chlorine atoms, and ulti- 
mately escape as chlorine gas. All cases of electrolysis are 
not as simple as this one, but it serves admirably as an intro- 
ductory illustration. It should be noted that the electric 
current does not tear the molecules apart, as was once sup- 
posed. The molecules of hydrochloric acid that dissociate 
are already dissociated before the electric current is intro- 
duced. The current upsets the electrical equilibrium be- 
tween the ions, so to speak, and they start at once on a mi- 
gration toward their proper electrodes where they lose their 
charges and become ordinary, uncharged atoms or atomic 
groups. Careful and extended experiments have not only 
demonstrated the actual movement of ions, but have deter- 



SOLUTION 143 

mined the rate of migration in many cases. Non-electrolytic 
solutions do not conduct electricity, because, according to 
the theory of electrolytic dissociation, they contain no ions. 
In this connection it is appropriate to emphasize the fact that 
only those solutions conduct electricity that have been found 
by other methods to contain ions. 

Electrolysis is often a complicated process, since the regen- 
erated atoms and atomic groups may interact chemically 
with the constituents of the solution and sometimes with 
each other or with the electrodes. The electrolysis of copper 
sulphate furnishes a typical illustration. The ions of a 
copper sulphate solution are copper ions (Cu ++ ) and sul- 
phate ions (S0 4 ). When this solution is electrolyzed, 
the copper ions (Cu ++ ) lose their electric charges at the 
cathode, become copper atoms (Cu), and adhere as metallic 
copper to the cathode. The sulphate ions (S0 4 ) lose 
their electric charges at the anode and become ordinary, 
uncharged atomic groups (S0 4 ). But this group of atoms is 
chemically unstable, and immediately interacts with the 
water around the anode, forming sulphuric acid (H 2 S0 4 ) and 
oxygen (0). The oxygen escapes, but the sulphuric acid 
mingles with the solution and dissociates into its ions. 
Similarly, a solution of sodium sulphate when undergoing 
electrolysis yields sulphuric acid at the anode and sodium 
hydroxide (NaOH) at the cathode; the electrolyte itself 
(sodium sulphate) furnishes directly only sodium and sul- 
phate ions, which lose their charges at the electrodes, and 
by their subsequent chemical interaction with the w T ater give 
the final result just stated. 

The so-called electrolysis of water is readily interpreted 
by the theory of electrolytic dissociation. Water itself does 
not conduct electricity to an extent which is comparable with 
the behavior of an electrolytic solution, because water dis- 
sociates only inappreciably and gives therefore an exceed- 



144 INORGANIC CHEMISTRY 

ingly small number of ions. A solution of sulphuric acid 
contains hydrogen ions (H + H + ) and sulphate ions (S0 4 ). 
When a current is passed through this solution, hydrogen 
ions migrate to the cathode, lose their electric charges, be- 
come hydrogen atoms, and eventually escape as hydrogen 
gas; the S0 4 -ions migrate to the anode, lose their electric 
charges, become S0 4 -groups, and interact with the water 
to form sulphuric acid and oxygen. The oxygen escapes as 
a gas, while the sulphuric acid dissociates into its ions. The 
water, therefore, is not split up directly into hydrogen and 
oxygen, as was formerly supposed. The two liberated 
gases are produced by the joint operations of electrolysis 
and subsequent chemical action, but the gases would not be 
liberated at all unless the ionization of the sulphuric acid 
had previously occurred in the solution. 

Electrolysis is a broad subject, and is not limited to aqueous 
solutions. In subsequent chapters frequent reference will 
be made to the electrolysis of molten substances, especially 
to the industrial applications which have become so impor- 
tant. Enough has been set forth at present, however, to 
show that the facts thus far revealed by electrolysis are in 
harmony with the theory of electrolytic dissociation. 

(d) Chemical Behavior of Electrolytic Solutions. — It has 
been pointed out that chemical action is often dependent upon 
the presence of water. Dry compounds like potassium 
chloride (KC1) and silver nitrate (AgN0 3 ) do not interact 
chemically, but if their solutions are mixed, a precipitate of 
silver chloride (AgCl) is immediately produced. On the 
other hand, there is no chemical action manifested when 
solutions of potassium chlorate (KC10 3 ) and silver nitrate are 
mixed, despite the fact that chlorine is a constituent of 
potassium chlorate. Furthermore, any chloride in solution 
will interact with silver nitrate in solution and produce a 
precipitate of silver chloride. These facts are typical of 



SOLUTION 145 

electrolytic solutions. Interpreted by the theory of elec- 
trolytic dissociation, they mean that reactions in solutions 
are due to some extent to ions. Dry or undissolved electro- 
lytes do not interact, because no ions are present, but in the 
case of dissolved electrolytes certain ions at once seek each 
other out in accordance with the fundamental principles of 
chemical action. If this action results in the formation of 
an insoluble compound, like silver chloride, this factor is 
removed from the scene of action as a precipitate and serves 
as visual evidence of the chemical change. Often ions are 
produced which cannot enter into chemical combination. 
Thus, a potassium chlorate solution contains potassium ions 
(K + ) and chlorate ions ((C10 3 ~), and when silver nitrate 
solution is added', the solution contains four kinds of ions, — 
potassium ions (K + ), chlorate ions (C10 3 ~), silver ions (Ag + ), 
and nitrate ions (N0 3 ~). But all compounds which might 
be formed by the various combinations of these ions are 
soluble. Hence the ions remain as such in the solution. It 
is for this reason that silver nitrate is effective in testing for 
hydrochloric acid or a soluble chloride, but not for other 
compounds containing chlorine, such as potassium chlorate 
(KCIO3) and chloroform (CHC1 3 ). Strictly speaking, the 
test is for chlorine ions or ionic chlorine, not for the element 
chlorine; and since the solutions of potassium chlorate and 
chloroform contain no ionic chlorine, the test fails with 
these compounds. Similarly, sulphuric acid and all soluble 
sulphates form insoluble barium sulphate (BaS0 4 ) when 
added to a solution of barium chloride (or any other soluble 
barium compound), because the sulphuric acid and sulphate 
solutions contain sulphate ions (S0 4 ~~), which combine 
with the barium ions (Ba ++ ) in the barium chloride solu- 
tion. But other sulphur compounds, such as sulphides, 
sulphites, and thiosulphates, do not form barium sulphate 
when added to barium chloride solution, because solutions 



146 INORGANIC CHEMISTRY 

of these compounds do not contain sulphate ions. It is clear 
from the above statements why a single test {i.e. the pre- 
cipitation of barium sulphate) is applicable to sulphuric 
acid and all soluble sulphates. All contain in solution a 
common ion (S0 4 ). 

Other properties besides the formation of precipitates 
are ascribed to ions and are often used as tests. Thus, the 
sour taste of all acids is attributed to hydrogen ions (H + ), 
which are common to acids. The color of solutions is also 
due to ions. Most ions are colorless, while solutions having 
a common colored ion have the same color. Thus, copper 
ions (Cu ++ ) are blue, and solutions containing such ions 
are blue, irrespective of the color of the undissolved copper 
compound. Cobalt ions (Co ++ ) are pink, and nickel ions 
(Ni ++ ) are green — colors usually exhibited by solutions of 
compounds of these elements. The migration of ions is often 
studied by means of colored ions. 

Common Ions. — Ions, as already stated, are electrically 
charged atoms or atomic groups. It is rather difficult for 
a beginner to determine what ions are present in a solution. 
The problem is simplified somewhat if the following general 
statements are borne in mind, (a) Hydrogen and metals 
form simple cations, (b) Non-metals (except hydrogen) form 
simple anions, (c) Some metals (e.g. Cr and Mn) and several 
non-metals (e.g. C, N, S, P) form compound ions — usually 
anions; e.g. HCO3-, N0 8 - HSOr, H 2 P0 4 ~, CrO", MnOr 
(note also OH~ and NH 4 + ). Certain elements likewise form 
complex ions; e.g. the silver-cyanogen ion ((Ag(CN) 2 )~), 
the silver-ammonia ion ((Ag(NH 3 ) 2 ) + ), and the copper- 
ammonia ion ((Cu(NH 3 ) 4 ) ++ ). The ions formed by the dis- 
sociation of the common compounds of the familiar ele- 
ments are shown in the following : — 



SOLUTION 



147 



Table of Common Ions 



Element 
or Group 


Ion 


Element 
or Group 


Ion 


Element 
or Group 


Ion 


Hydrogen . . 


H+ 


Calcium . . 


Ca++ 


Aluminium . 


A1+++ 


Sodium . . . 
Potassium . . 


Na+ 


Barium . . . 
Copper . . . 


Ba++ 
Cu++ 


Antimony . 
Bismuth . . 


Sb+++ 
Bi+++ 


Silver .... 


Ag+ 


Zinc .... 


Z11++ 


Iron (ic) . . 


Fe+++ 


Ammonium . 


NH 4 + 


Magnesium . 


Mg++ 


Tin (ic) . . . 


Sn++++ 


Mercury (ous) 
Chlorine . . . 


Hg+ 

ci- 


Lead .... 
Iron (ous) . 


Pb++ 

Fe++ 






Bromine . . . 
Iodine .... 


Bi~ 
I- 


Mercury (ic) 
Tin (ous) . 


Hg++ 
Sn++ 






Nitrate .... 
Chlorate . . . 
Hydroxyl . . 


N0 3 - 

CIO3- 

OH- 


Sulphate . . 
Sulphide . . 
Carbonate . 
Chromate . 


S0 4 ~ 

s — 

C0 3 — 
Cr0 4 — 










Di chromate 


Cr 2 7 — 







Many deductions arising from this rather compact table will 
be considered in the succeeding pages. 



Summary. — The salient points discussed in this chapter 
may be summarized as follows : The properties of a solution 
are mainly dependent upon the solute and its condition in 
the solvent. Such properties as osmotic pressure, freezing 
point, and boiling point, are influenced by the number of 
independent particles present in the solution. When these 
three properties are measured independently in a given solu- 
tion, the values agree. But when the values are compared 
under parallel conditions of measurement, it is found that 
electrolytes in solution yield a larger number of independent, 
individual particles than non-electrolytes; i.e. in electro- 
lytic solutions some of the molecules of the electrolytes disso- 
ciate into ions. These ions are electrically charged atoms 
or atomic groups. Physically they act much like molecules. 



148 INORGANIC CHEMISTRY 

Furthermore, solutions of electrolytes differ from solutions 
of non-electrolytes in conducting electricity and exhibiting 
marked chemical activity. A study of these two character- 
istics confirms the assumption of the existence of electrically 
charged and chemically active particles in the solution. 

Conclusion. — The theory of electrolytic dissociation as 
outlined in the foregoing pages is not an adequate explana- 
tion of all the facts of solution. It applies chiefly to dilute 
aqueous solutions of three classes of substances. Doubtless 
the present form of the theory will sometime be modified to 
cover certain facts not at present within its scope. 

Exercises 

1. Write out the formulas of the ions formed when the following 
compounds are dissolved separately in considerable water: Potassium 
chloride, silver nitrate, sodium chlorate, ammonium sulphate, copper 
nitrate, calcium chloride, zinc sulphate, potassium dichromate, 
calcium hydroxide. 

2. Write the equations for the following by applying the method 
for making equations outlined in Chapter VII: (a) Iron and sulphur 
combine in the ratio of 7 to 4. (b) Ammonia gas and hydrochloric 
acid gas form ammonium chloride, (c) Magnesium and hydrochloric 
acid form hydrogen and magnesium chloride. 

3. Discuss: (a) Electrolytes depress the freezing point abnor- 
mally ; (b) ions migrate to their respective electrodes. 

4. Write the following as ionic equations : (a) Potassium sul- 
phate and barium chloride form barium sulphate and potassium 
chloride; (6) sodium bromide and silver sulphate form silver bro- 
mide and sodium sulphate. 



CHAPTER X 
Acids, Bases, and Salts — Neutralization 

Introduction. — Many chemical compounds fall naturally 
into one of three groups, long known as acids, bases, and salts. 
Each group has its characteristic properties, though the 
groups are closely related and sometimes overlap. Many 
familiar substances belong to these groups. 

Acids. — The common acids are sulphuric acid (H 2 S0 4 ), 
hydrochloric acid (HC1), nitric acid (HX0 3 ), and acetic acid 
(C 2 H 4 2 ). Many acids are liquid, as sulphuric and nitric; 
a few are gases, as hydrochloric; others are solid, as tartaric, 
citric, oxalic. Most acids are rather soluble in water, and 
such solutions are popularly called acids. These solutions 
may be dilute or concentrated, and the general properties 
vary somewhat with the concentration. Concentrated acids 
are usually corrosive and should be handled with caution, 
even when one is familiar with their properties. 

Many familiar substances are acids or contain them. 
Vinegar, pickles, and similar relishes contain dilute acetic 
acid. Lemon juice is mainly citric acid. Sour milk con- 
tains lactic acid. Unripe fruits, sour bread, and sour w T ines 
contain acids. "Soda water" is a solution of carbonic acid 
(or more accurately carbon dioxide), and "acid phosphate" 
is a solution of a sour calcium phosphate. 

Properties of Acids. — (1) Acids, if dissolved in water, 
usually have a sour taste. The early chemists detected this 
fact, and the term acid (from the Latin word acidus, 

149 



150 INORGANIC CHEMISTRY 

sour) emphasizes this property. (2) Solutions of acids 
redden the coloring matter called litmus. Solutions which 
act thus on blue litmus are described as acid, as containing 
an acid, or as having an acid reaction. (3) Most acids 
liberate free hydrogen gas when their solutions interact with 
metals. (4) Solutions of acids conduct electricity. 

Composition of Acids. — All acids contain hydrogen, 
which is liberated in the free state when certain metals and 
acids interact. Most acids contain oxygen. For many 
years it was thought that oxygen was an essential component 
of all acids, and the name oxygen (derived from Greek 
words meaning "acid producer") was given to this element 
by Lavoisier because of this belief. (See Discovery of Oxygen.) 
We know now that hydrogen, not oxygen, is the essential 
constituent of all acids. Another necessary constituent of 
acids is a non-rnetallic element like nitrogen or sulphur. 
For this reason it is sometimes convenient to think of non- 
metals as the elements which form acids. Thus, sulphuric 
acid contains sulphur, besides hydrogen and oxygen; while 
hydrochloric acid contains only chlorine, besides hydrogen. 
The important non-metals which form familiar acids are 
boron, carbon, silicon, nitrogen, phosphorus, sulphur, flu- 
orine, chlorine, bromine, and iodine. 

Definition of an Acid. — For many years an acid was 
defined as a compound producing a sour solution which 
reddens blue litmus, or as a compound which interacts 
chemically with a base, thereby forming a salt, or as a 
compound containing hydrogen which can be replaced by a 
metal. These definitions emphasize certain properties of 
acids, but they are not inclusive. According to the theory 
of electrolytic dissociation, an acid is a compound whose 
solution contains hydrogen ions (H + ). The sour taste, 
behavior toward litmus, and liberation of hydrogen are due 



ACIDS, BASES, AND SALTS 151 

to the hydrogen ions which are common to all solutions of 
acids. 

Bases. — - The common bases are sodium hydroxide 
(NaOH), potassium hydroxide (KOH), ammonium hydrox- 
ide (NH 4 OH), and calcium hydroxide (Ca(OH) 2 ). They are 
soluble in water, and such solutions are called bases; solu- 
tions of the very soluble bases (sodium hydroxide and po- 
tassium hydroxide) are often called alkalies. Alkalies, like 
concentrated acids, are corrosive, and should be handled 
carefully. Concentrated solutions of sodium and potas- 
sium hydroxides are very corrosive, and for this reason are 
called caustic alkalies (from the Latin causticus, burning). 

Bases are components of familiar substances. Thus, 
ammonia is a solution of ammonium hydroxide. Lime- 
water and baryta water are solutions of the sparingly soluble 
bases calcium hydroxide and barium hydroxide (Ba(OH) 2 ) 
respectively. Lye is a concentrated solution of sodium 
hydroxide or potassium hydroxide (or both). 

Properties of Bases. — (1) Strong, soluble bases have a 
bitter, often biting, taste; many, especially the very soluble 
ones, have a slippery feeling. (2) Soluble bases turn red 
litmus blue. Substances which act thus on red litmus are 
described as basic or alkaline, as having an alkaline reaction, 
or as containing a base. (3) Solutions of bases conduct 
electricity. 

Composition of Bases. — All bases contain hydrogen and 
oxygen. They also contain a metal, such as sodium. The 
hydrogen and oxygen are the invariable constituents. But 
it is often convenient to regard metals as the elements which 
form bases, just as the non-metals form acids. Thus, the 
base sodium hydroxide is a compound of the metal sodium 
with hydrogen and oxygen. The important metals which 



152 INORGANIC CHEMISTRY 

form familiar bases are sodium, potassium, calcium, and 
barium. 

Definition of a Base. — A base was formerly defined as 
any compound which has a bitter taste, turns red litmus blue, 
and interacts chemically with an acid, thereby forming a 
salt. This definition emphasizes certain properties of a 
base, but it is defective. According to the theory of elec- 
trolytic dissociation, a base is a compound whose solution 
contains hydroxyl ions (OH"). The metal is not the con- 
stituent which gives a base its characteristic properties. 
These are due to the hydroxyl ions which are common to all 
solutions of bases. 

Salts. — This is a large and varied class of compounds. 
The most familiar member is sodium chloride (NaCl). It 
is common salt or table salt, and has been known for ages. 
Doubtless this class of chemical compounds received its 
name from the general resemblance many of them bear to 
common salt. Most salts are solids and are soluble in 
water, although the solubility varies between wide limits. 

Properties of Salts. — (1) Salts often have the well-known 
salty taste, though some are bitter, others are astringent, and 
a few have no characteristic taste. Certain salts have a 
sour taste. (2) Salts do not act uniformly on litmus. 
Some turn the red to blue, others turn the blue to red, and 
many have no action whatever on litmus. Those salts 
whose solutions do not respond to the litmus test are said to 
be neutral or to have a neutral reaction. This indifference 
to litmus is not a decisive test for a salt, since many other 
substances, water for example, have no action on litmus. 
The term neutral is applied to substances which do not 
change the color of litmus, whether or not they are salts or 
contain salts. (3) Solutions of salts conduct electricity. 



acids; bases, and salts 153 

Composition of Salts. — Salts contain invariably a metal 
and a non-metal (which is not hydrogen or oxygen). Most 
salts also contain oxygen. Thus, potassium nitrate contains 
the metal potassium and the non-metal nitrogen, besides 
oxygen; while potassium chloride contains potassium and 
the non-metal chlorine, but no oxygen. A few salts contain 
hydrogen besides the characteristic metal and non-metal. 
Thus,' sodium bicarbonate contains hydrogen besides the 
metal sodium, the non-metal carbon, and the non-metal 
oxygen. These rather general statements indicate the great 
variety of salts. Salts will soon be further discussed. 

Neutralization. — The nature and interrelation of acids, 
bases, and salts are shown clearly by their chemical relations. 
When a solution of an acid and a base are mixed in the 
proper proportion, they interact completely. The final 
solution has none of the properties of an acid or a base, but 
it has the properties characteristic of a salt. That is, the 
acid and base destroy more or less completely the marked 
properties of each other and produce a salt, the latter being a 
compound which has few, if any, of the properties of the 
original acid and base. The acid and base neutralize each 
other. An illustration will make this point clear. When 
hydrochloric acid and sodium hydroxide interact, sodium 
chloride and water are formed. The chemical change can 
be written thus : — 

HC1 + NaOH = NaCl + H 2 

Hydrochloric Sodium Sodium Water 

Acid Hydroxide Chloride 

This equation represents the facts which have been repeatedly 
verified by experiment. The chemical change in which an 
acid and a base neutralize each other and form a salt and 
water is called neutralization. Taking this equation as a 
type of the chemical change which occurs in neutralization, 



154 INORGANIC CHEMISTRY 

it is clear that (1) the metal of the base takes the place of 
the hydrogen of the acid, thereby forming a salt, while 
(2) the hydrogen of the acid combines with the hydrogen 
and oxygen of the base to form water. In neutralization 
the hydrogen and oxygen of the base act as a unit. This 
group of atoms (OH), as already stated, is called hydroxyl. 
Hydroxyl does not exist free and uncombined like elements 
and compounds, but it acts like a single atom in many 
chemical changes. It is called a radical. 

Neutralization illustrates double decomposition. In the 
chemical change just cited both the hydrochloric acid and 
the sodium hydroxide are decomposed and their parts are 
recombined in a different way; i.e. sodium chloride and water 
are the new compounds resulting from the recombination. 

Neutralization when interpreted by the theory of elec- 
trolytic dissociation is really the union of hydrogen ions with 
hydroxyl ions. Suppose solutions of hydrochloric acid and 
sodium hydroxide are mixed in the proper proportions. The 
mixture at first contains ions of hydrogen, chlorine, sodium, 
and hydroxyl. But the ions of hydrogen and of hydroxyl 
immediately unite to form molecules of water, because water 
does not dissociate into ions to any appreciable extent. The 
final solution is neutral, because it contains only ions of 
sodium and chlorine, the acid ions (H + ) and the basic ions 
(OH") having been removed by their combination into 
molecules of water. The equation for the mutual neutraliza- 
tion of hydrochloric acid and sodium hydroxide might be 
written as an ionic equation, thus : — 

H+ + CI- + Na+ + OH" ->C1- + Na + + H 2 

The ionic equations for the mutual neutralization of other 
pairs of acids and bases are similarly written. In the case 
just described, the ions of sodium and of chlorine remain 
uncombined until the solution is evaporated; but as the con- 



ACIDS, BASES, AND SALTS 



155 



centration increases, the ions unite and form molecules of 
sodium chloride. The latter, as already stated, is one type 
of the varied class of compounds called salts. The salts 
resulting from the combination of ions can be obtained as 
solids by the usual processes of evaporation to dryness or by 
crystallization. It is evident, therefore, that neutralization 
in a broad sense is the mutual destruction of an acid and a 
base which results in the formation of a salt and water. In 
a narrow sense it is the formation of molecules of water from 
the hydrogen ions of the acid and the hydroxyl ions of the 
base. This latter interpretation is supported by experi- 
mental evidence. Heat is liberated when these ions unite to 
form water. If neutralization is merely the combination of 
ions of hydrogen with ions of hydroxyl, 
then the same amount of heat should be 
liberated when a given weight of water 
is formed, whether the ions come from 
hydrochloric acid and sodium hydroxide 
or from any other 
. pair of acid and 
base. Experi- 
ment shows that 
the heat of neu- 
tralization, as it 
is called, is the 
same in all cases 
of neutralization, 
when the solu- 
tions are dilute 
and other thermal 
changes do not 
occur. It is expressed in terms of a unit called the calorie, and 
when 18 gm. of water are formed by the act of neutralization, 
13,700 calories are liberated. (See Calorie, Chapter XL) 



11— 



h^d: 



12- 



— J 



.JI 



-III 




Fig. 18. — Burettes. Enlarged section (on the left) 
shows graduations and curved surface of the solu- 
tion, called the meniscus. Correct reading is along 
line I. 



156 INORGANIC CHEMISTRY 

Neutralization is frequently brought about by using bu- 
rettes (Fig. 18). These are graduated tubes provided with 
a stopcock to regulate the exit of the solution. When prop- 
erly filled, one with an acid {e.g. sulphuric) and the other with 
a base (e.g. sodium hydroxide), a measured volume of the 
acid is allowed to flow into a beaker and a few drops of lit- 
mus solution are added. The solution of course turns red. 
The sodium hydroxide solution is allowed to drop in slowly 
and the mixture is stirred with a glass rod. As long as an 
excess of acid is present, the color remains red. After 
a time, however, the color becomes purple, and an ad- 
ditional drop of sodium hydroxide turns the solution blue, 
showing that the acid has been neutralized by the base. 
The volume of sodium hydroxide used is noted. If solutions 
of known strength are used, then the weights (found from the 
concentration of the solutions) used will be in the same ratio 
as the weights in the equation: — 

H 2 S0 4 + 2. NaOH = Na 2 S0 4 + 2 H 2 

Sulphuric Sodium Sodium Water 

Acid Hydroxide Sulphate 36 

98 80 142 

If the concentration of the sodium hydroxide solution is 
unknown, it can be found by the proper proportion, because 
the weights involved in the chemical change are always in the 
ratio given in the corresponding equation. Neutralization 
when conducted by means of accurate apparatus and certain 
solutions of known strength is an efficient method of quan- 
titative analysis and is one of a class called volumetric 
methods. 

Classification of Salts. — It will be recalled that salts have 
no distinctive class property like acids and bases, such as 
the taste and behavior with litmus. From the standpoint 
of the chemical change which occurs in complete neutral- 



ACIDS, BASES, AND SALTS 157 

ization and in analogous cases where the chemical change is 
not complete, a salt is a compound formed (1) by the substi- 
tution of a metal for all or part of the hydrogen of an acid, 
or (2) by the substitution of a non-metal (or non-metallic 
group like S0 4 or N0 3 ) for all or some of the hydroxyl groups 
of a base. There are, therefore, three classes of salts, — 
normal, acid, and basic. They are prepared in various ways, 
but it is convenient to regard them as having been produced 
from an acid or a base by substitution. Salts formed by re- 
placing all the hydrogen of an acid by a metal are called 
normal salts, e.g. sodium sulphate, Na 2 S0 4 . On the other 
hand, salts formed by replacing only part of the hydrogen 
of an acid by a metal are called acid salts. Thus, acid sodium- 
sulphate (HNaS0 4 ) may be regarded as derived from sul- 
phuric acid by replacing only one of the atoms of hydrogen 
by one atom of sodium, though of course the salt is not pre- 
pared in this way. - -. . 

Only the acids containing two or more replaceable atoms 
of hydrogen can, as a rule, form acid salts ; e.g. sulphuric 
acid (H 2 S0 4 ) and phosphoric acid (H3PO4). Acids are often 
classified by the number of their hydrogen atoms which can 
be replaced by a metal. This varying power of replaceability 
is called basicity. A monobasic acid contains only one atom 
of replaceable hydrogen in a molecule ; e.g. nitric acid, HN0 3 . 
A molecule of acetic acid (C 2 H 4 2 ) contains four atoms of hy- 
drogen, but for reasons which are too complex to state here, 
only one of these atoms can be replaced by a metal ; it is there- 
fore monobasic. Dibasic and tribasic acids are those that con- 
tain respectively two and three replaceable hydrogen atoms ; 
e.g. sulphuric acid (H 2 S0 4 ) and phosphoric acid (H 3 P0 4 ). 
Normal salts may also be regarded as formed by the replace- 
ment of all the hydroxyl groups of a base by non-metallic 
atoms or atomic groups. Thus, bismuth nitrate (Bi(N0 3 )3) 
is a normal salt and may be regarded as formed by the sub- 



158 



INORGANIC CHEMISTRY 



stitution of three N0 3 -groups for the three hydroxyl groups 
in bismuth hydroxide (Bi(OH) 3 ), while basic bismuth nitrate 
(Bi(OH) 2 N0 3 ) — a basic salt — : may be regarded as formed 
by the substitution of one N0 3 -group for one of the three 
OH-groups in the hydroxide. 

Bases, like acids, are classified according to their varying 
power of replaceability. This power is called acidity. Bases 
are called monacid, diacid, triacid bases, etc., according to the 
number of the replaceable hydroxyl groups present in a mole- 
cule. Thus, sodium hydroxide (NaOH) is a monacid base, 
calcium hydroxide (Ca(OH) 2 ) is a diacid base, and aluminium 
hydroxide (Al(OH) 3 ) is a triacid base. Only bases having 
two or more replaceable hydroxyl groups form basic salts. 

The relations of acids, bases, and salts may be represented 
by the following scheme : — 



Acid 
H 2 S0 4 

Sulphuric Acid 



Base 

Zn(OH) 2 

Zinc Hydroxide 



^Normal Salts" 

Na 2 S0 4 

Sodium Sulphate 

ZnCl 2 

Zinc Chloride 



Acid Salt 
HNaS0 4 

Acid Sodium Sulphate 



Basic Salt 
Zn(OH)Cl 

Basic Zinc Chloride 



Preparation of Salts. — It must not be concluded from 
the foregoing discussion of the kinds of salts that they are 
always prepared in the laboratory by mixing acids with 
bases or metals. It is only necessary to provide the metallic 



ACIDS, BASES, AND SALTS 159 

or non-metallic constituent chemically, so to speak. Salts 
can be prepared in several ways. The interaction of an acid 
and a base has been mentioned. The interaction of acids 
with oxides of certain metals produces salts. Sodium oxide 
and sulphuric acid interact and form the salt sodium sulphate, 
thus: — 

Na 2 + H 2 S0 4 = Na 2 S0 4 + H 2 

Sodium Sulphuric Sodium Water 

Oxide Acid Sulphate 

A metal and an acid act similarly. Zinc and sulphuric acid, 
as already stated, form the salt zinc sulphate, thus: — 



Zn + H 2 S0 4 = 


= ZnS0 4 


+ 2H 


Zinc Sulphuric 
Acid 


Zinc 
Sulphate 


Hydrogen 



Carbonates interact with acids and form other salts. Cal- 
cium carbonate and hydrochloric acid form the salt calcium 
chloride, thus: — 



CaC0 3 


+ 


2HC1 


= CaCl 2 


+ 


co 2 


+ H 2 


Calcium 




Hydrochloric 


Calcium 




Carbon 


Water 


Carbonate 




Acid 


Chloride 




Dioxide 





Sometimes two salts interact in solution and form other 
salts by double decomposition. Sodium chloride and silver 
nitrate form the salts silver chloride and sodium nitrate, 
thus: — 



NaCl 


+ 


AgN0 3 = 


= AgCl 


+ 


NaN0 3 


Sodium 




Silver 


Silver 




Sodium 


Chloride 




Nitrate 


Chloride 




Nitrate 



Salts interact with certain acids and form salts and other 
acids. Sodium chloride and sulphuric acid form hydro- 
chloric acid and the salt sodium sulphate, thus: — 

2 NaCl + H 2 S0 4 = 2 HC1 + Na 2 S0 4 

Sodium Sulphuric Hydrochloric Sodium 

Chloride Acid " Acid Sulphate 



160 



INORGANIC CHEMISTRY 



Nomenclature of Acids. — Oxygen is a component of most 
acids, and the names of these acids correspond to the pro- 
portion of oxygen which they contain. The best known 
acid of an element usually has the suffix -ic; e.g. sulphuric 
(H 2 S0 4 ), nitric (HN0 3 ), phosphoric (H 3 P0 4 ). If an element 
forms another acid, containing less oxygen, this acid has the 
suffix -ous; e.g. sulphurous (H 2 S0 3 ), nitrous (HN0 2 ), phos- 
phorous (H3PO3). Some elements form an acid containing 
less oxygen than the -ous acid; these acids retain the suffix 
-ous , and have in addition the prefix hypo-; e.g. hyposulphu- 
rous (HS0 2 ) , hypophosphorous (H 3 P0 2 ) , hypochlorous (HCIO) . 
Hypo- means under or lesser. If an element forms an acid 
containing more oxygen than the -ic acid, such an acid re- 
tains the suffix -ic, and has in addition the prefix per-; e.g. 
persulphuric (H 2 S 2 8 ), perchloric (HC10 4 ). The prefix per- 
means beyond or over. The few acids which contain no 
oxygen have the prefix hydro- and the suffix -ic; e.g. hydro- 
chloric (HCl), hydrobromic (HBr), hydrofluoric (HF), and 
hydriodic (HI). It should be noticed that the suffixes ic- 
and -ous are not always added to the whole name of the 
element, but often to some modification of it. 

The nomenclature of acids is well illustrated by the series 
of chlorine acids: — 

Acids of the Element Chlorine 



Name 


Formula 


Hydrochloric 

Hypochlorous 

Chlorous 


HCl 

HCIO 

HCIO2 


Chloric 

Perchloric 


HCIO3 
HCIO4 



Not all elements form a complete series of acids, but the 
nomenclature usually agrees with the above principles. 



ACIDS, BASES, AND SALTS 161 

Some acids have commercial names. Thus, sulphuric 
acid is often called oil of vitriol, and hydrochloric acid is 
known as muriatic acid. Acids in which carbon is an essen- 
tial constituent end in -ic, but they are often arbitrarily 
named. (See Organic Acids.) 

Nomenclature of Bases. — There is no general rule for the 
nomenclature of bases, as in the case of acids. Since most 
bases contain hydrogen and oxygen, they are called hydrox- 
ides. The term hydrate is sometimes used as a synonym 
of hydroxide, but it more correctly describes those com- 
pounds in which water is one constituent. Thus, crystal- 
lized salts containing water of ciystallization are often 
called hydrates. The term alkali emphasizes general prop- 
erties rather than suggests specific composition and is now 
applied to the very active bases such as sodium and potas- 
sium hydroxides. Hydroxides are distinguished from each 
other by placing the name of the metal before the word 
hydroxide; e.g. sodium hydroxide, potassium hydroxide, 
calcium hydroxide. 

The common hydroxides have long been known by special 
names. Thus, a solution of calcium hydroxide is sometimes 
called limew r ater. Ammonium hydroxide solution is some- 
times called ammonia water or simply ammonia, and it was 
formerly called volatile alkali. The hydroxides of sodium 
and potassium are often called caustic soda and caustic 
potash, and occasionally the term fixed alkali is used to 
emphasize the fact that they are non-volatile. 

Nomenclature of Salts. — The names of salts containing 
oxygen are derived from the name of the corresponding acid. 
The characteristic suffix of the acid is changed to indicate 
this relation. Thus, the suffix -ic becomes -ate, and the suffix 
-ous becomes -ite. Hence : — 



162 INORGANIC CHEMISTRY 

Sulphuric acid forms sulphates. 
Sulphurous acid forms sulphites. 
Nitric acid forms nitrates. 
Nitrous acid forms nitrites. 
Chloric acid forms chlorates. 
Hypochlorous acid forms hypochlorites. 
Permanganic acid forms permanganates. 

The name of the replacing metal is retained ; e.g. potassium 
chlorate, sodium sulphate, calcium hypochlorite, potassium 
permanganate. Notice that the prefixes hypo- and per- are 
not changed. 

The names of salts containing only two elements, follow- 
ing the general rule for binary compounds, end in -ide. This 
suffix is added to a modification of the name of the non-metal, 
giving the names chloride, bromide, sulphide, fluoride, etc. 
The prefix hydro- which is contained in the name of the acid 
is omitted. Thus, the sodium salt of hydrochloric acid is 
sodium chloride ; similarly, there are the names potassium 

chloride, calcium fluoride, and sodium iodide. 
4- 

Relation of Oxides to Acids and Bases. — Most non- 
metallic elements form oxides which unite with water and 
produce an acid. The oxides of many metallic elements, on 
the other hand, unite with water and produce bases. The 
two oxides of the non-metal sulphur act thus : — 



so 2 


+ 


H 2 


= H 2 S0 3 


Sulphur 
Dioxide 




Water 


Sulphurous 
Acid 


so 3 


+ 


h 2 o 


= H 2 S0 4 


Sulphur 
Trioxide 




Water 


Sulphuric 
Acid 


The oxide of the metal calcium acts thus : — 


CaO 


+ 


H 2 


= Ca(OH) 2 


Calcium 
Oxide 




Water 


Calcium 
Hydroxide 



ACIDS, BASES, AND SALTS 163 

Oxides of non-metals which unite with water and produce 
acids are called acid forming oxides or acid anhydrides, i.e. 
literally, substances without water. Examples are carbonic 
anhydride (C0 2 ), sulphuric anhydride (S0 3 ), phosphoric 
anhydride (P 2 5 ). Oxides of metals which produce basic 
hydroxides are base forming oxides or basic anhydrides. 
Examples are calcium oxide (CaO), sodium oxide (Na 2 0), 
barium oxide (BaO) . A few oxides behave exceptionally. It 
is convenient to regard an acid anhydride as the root or basis 
of its corresponding acid, and a basic anhydride as the root 
of its hydroxide. 

The fact that many non-metallic oxides redden moist blue litmus 
led Lavoisier into the erroneous belief that oxygen is an essential 
constituent of all acids. And some authorities even now speak 
(incorrectly) of these oxides as acids; thus, carbon dioxide (CO2) is 
occasionally called carbonic acid. The compounds which Lavoisier 
called acids were anhydrides. And it was not until about 1811 that 
Davy showed (1) that some acids do not contain oxygen (e.g. hydro- 
chloric acid, HC1), and (2) that the so-called acids of Lavoisier are 
not real acids until they obtain hydrogen from the water in which 
they dissolve. 

Degree of Dissociation of Acids, Bases, and Salts. — The 
degree to which acids, bases, and salts dissociate is due to 
two factors, viz. the nature of the substance itself and the 
concentration of the solution. In general, dissociation is 
slight in a concentrated solution, and increases as the solu- 
tion becomes more and more dilute. Thus, in concentrated 
nitric acid (62 per cent) the per cent of ionized substance is 
only about 9.6, while in a dilute solution (.63 per cent) the 
per cent is about 90. The different degrees of dissociation 
of acids, bases, and salts may be readily seen by consulting 
the table of the dissociation of these substances on the next 
page. In order to compare the varying degrees of dissocia- 
tion, solutions of the same relative strength (at a fixed tem- 
perature) must be- taken. Usually normal solutions at 18° C. 



164 



INORGANIC CHEMISTRY 



are selected, i.e. solutions which contain in a liter a number 
of grams numerically equal to the molecular weight divided 
by the number of replaceable hydrogen atoms present (or 
their metallic equivalent). For example, a normal solution 
of hydrochloric acid contains 36.5 gm. to the liter, i.e. (35.5 
+ l)-f- 1; but since sulphuric acid contains two replace- 
able hydrogen atoms, its normal solution contains (2 + 32 
+ 64) -T- 2 = 49 gm. in each liter of solution. Similarly, a 
normal solution of sodium hydroxide contains (23 + 16 + 1) 
-f- 1 = 40 gm., while one of potassium sulphate contains 
(78 + 32 + 64) -r- 2 = 87 gm. to the liter of solution. Nor- 
mal solutions are often designated by the letter N, and differ- 
ent concentrations are indicated correspondingly, e.g. tenth 

normal by — , twice normal by 2 N. The approximate 

per cent of dissociation of certain acids, bases, and salts is 
tabulated below; the solutions are one tenth normal (at 18°C). 



Table of Dissociation of Acids, Bases, and Salts 


Electrolyte 


Per Cent of Dissociation 


Nitric acid (H+, N0 3 -) 


90 


Hydrochloric acid (H+, CI - ) • . 


90 


Sulphuric acid (H+, HS0 4 -) 


60 


Acetic acid (H+, C 2 H 3 2 ~) . . '. ... . 


1.3 


Carbonic acid (H+, HC0 3 -) ....... 


.17 


Hydrocyanic acid (H+, CN~) 


.01 


Potassium hydroxide (K+, OH - ) 


90 


Sodium hydroxide (Na+, OH - ) 


90 


Ammonium hydroxide (NH 4 +, OH-) .... 


1.4 


Potassium chloride (K+, CI - ) 


86 


Sodium chloride (Na+, CI - ) 


84 


Potassium nitrate (K+, N0 3 ~) 


83 


Potassium sulphate (2 K+, S0 4 ~ ~) . . • • 


72 


Silver nitrate ((Ag+, N0 3 ~) ....... 


86 


Copper sulphate (Cu++, S0 4 ~ -) 


37 



ACIDS, BASES, AND SALTS 165 

An examination of the table above shows more or less 
equality of dissociation among salts, but a rather wide 
variation in the case of both acids and bases. The term 
strong is applied to those acids and bases which dissociate 
to a marked degree, and the term weak to those whose dis- 
sociation is limited. 

Salts and the Theory of Electrolytic Dissociation. — Ac- 
cording to the theory of electrolytic dissociation, solutions of 
normal salts contain neither hydrogen nor hydroxy 1 ions; 
and, as a rule, they are neutral to litmus. But solutions of 
acid salts contain the ions characteristic of an acid (i.e. 
H-ions or H + ) as well as the ions of a salt. While solutions 
of basic salts contain the ions characteristic of a base (i.e. 
hydroxyl ions or OH~) as well as the ions of a salt. A salt 
may be defined in terms of the theory of electrolytic disso- 
ciation as the compound ultimately formed by the union of 
one or more metallic ions of the base and one or more non- 
metallic ions of an acid, supplemented in some cases by one 
or more hydrogen or hydroxyl ions. In other words, salts 
are chemical compounds which have a certain composition 
but not necessarily uniform behavior. For example, take 
the behavior toward litmus. Normal salts are neutral, acid, 
or basic; acid salts are acid or nearly neutral; most basic 
salts are nearly insoluble in water, and exhibit a faint 
reaction toward litmus. The terms normal, acid, and 
basic as applied to salts indicate their composition but 
do not describe their properties. An explanation of this 
varying behavior of salts toward litmus is offered by the 
theoiy of electrolytic dissociation. Hitherto, pure water 
has been referred to as a non-electrolyte, and in most solu- 
tions there is no evidence of its dissociation into the ions 
H + and OH"". In some cases, however, the very slight 
ionization of water becomes a significant factor in establish- 



166 INORGANIC CHEMISTRY 

ing the properties revealed by the solution; i.e. its ions, 
though comparatively very few in number, interact with the 
ions of certain substances and thereby produce very interest- 
ing results. Thus, sodium carbonate yields the ions 2 Na + 
and C0 3 ~~; but C0 3 -ions are unstable and combine with 
H-ions to form the ion HC0 3 ~ — the last named being one 
of the ions yielded by the slightly dissociated acid H 2 C0 3 . 
The removal of H-ions leaves an excess of OH-ions, which 
give the solution an alkaline reaction toward litmus. On 
the other hand, copper sulphate (a normal salt from the 
standpoint of composition) has an acid reaction. It yields 
the ions Cu ++ and S0 4 ~ ~ ; but the Cu-ions combine with 
OH-ions to form copper hydroxide (Cu(OH) 2 ), which dis- 
sociates only to a very slight extent. The removal of OH-ions 
leaves an excess of H-ions in the solution and gives it an 
acid reaction. For similar reasons the litmus reaction of 
potassium carbonate and potassium cyanide is alkaline, 
while the reaction of ferric chloride is acid. 

Acid and basic salts are readily interpreted by the theory. 
Acid sodium sulphate (HNaS0 4 ) yields Na + and HS0 4 ", 
but the latter ion dissociates to some extent into H + and 
S0 4 ~ ~ ; the solution is therefore rendered acid to litmus by 
the, H-ions. But acid sodium carbonate (HNaC0 3 ) is nearly 
neutral to litmus. It yields the ions Na + and HC0 3 " ; but 
since the latter dissociates to only a very slight extent, the 
solution is not provided with an excess of either H-ions or 
OH-ions, and therefore ' has only a fairly neutral reaction. 
Basic salts may be similarly explained. 

Chemical changes like those cited in the first paragraph are 
examples of hydrolysis, i.e. a chemical change involving 
water and certain salts. It is typically exhibited by salts 
derived from strong bases {e.g. NaOH and KOH) and weak 
acids {e.g. H 2 C0 3 and HCN), or from weak bases {e.g. 
Cu(OH) 2 and Fe(OH) 3 ) and strong acids {e.g. HC1 and 



ACIDS, BASES, AND SALTS 167 

H 2 S0 4 ). The behavior of the final solution toward litmus 
depends upon the composition of the salt, as the following 
equations show. The ordinary chemical equation for the 
hydrolysis of sodium carbonate is — 

Na 2 C0 3 + H 2 = HNaC0 3 + NaOH 

Sodium Water Acid Sodium Sodium 

Carbonate Carbonate Hydroxide 

The ionic equation is — 

2 Na + + CO3- - + H + + OH- -^2 Na + + HC0 3 - + OH" 

The corresponding equations for the hydrolysis of copper 
sulphate are — ■ 



CuS0 4 + 2 H 2 = Cu(OH) 2 + H 2 S0 4 

Copper Water Copper Sulphuric 

Sulphate Hydroxide Acid 

Cu ++ + S(V - + 2 H + + 2 OH--^Cu(OH) 2 + SOr " + 2 H+ 



Problems and Exercises 

1. Name the hydroxide corresponding to sodium, potassium, 
calcium, barium, zinc, lead, copper. 

2. Name the potassium salt of manganic acid, calcium salt of 
hydrofluoric acid, sodium salt of carbonic acid, potassium salt of 
tartaric acid, lead salt of chromic acid, potassium salt of hydrobromic 
acid, potassium salt of permanganic acid. 

3. Name the sodium salt of hydrochloric acid. Name the cor- 
responding salt of potassium, lead, calcium, barium, zinc, silver. 

4. Name the same salts of nitric acid. Of nitrous acid. 

5. Name the same salts of sulphuric acid. Of hypochlorous acid. 
Of perchloric acid. 

6. Give the name and formula of the ions, if any, formed by 
(a) potassium sulphate, (b) silver chloride, (c) barium sulphate, 
(d) lead nitrate, (e) lead sulphate, (/) lead chromate. 

7. Classify into acids, bases, salts, and anhydrides : S0 3 , Pb (0H) 2 , 
HBr, KCIO4, C0 2 , H3PO4, P2O5, NH4CI, Ag 2 S0 4 , Ca(N0 3 ) 2 , HN0 2 , 
Ba(OH) 2 , HI, HNaC0 3 , Cu(OH) 2 , FeCl 3 , HNaS0 4 , NaN0 2 , S0 2j CaO. 



168 INORGANIC CHEMISTRY 

8. What is the percent of hydrogen in (a) sulphuric acid, (6) hy- 
drochloric acid, (c) nitric acid ? 

9. What is the percent of OH in (a) NaOH, (b) KOH, (c) NH 4 - 
OH, (d) Ca(OH) 2 ? 

10. Write equations for the following reactions : (a) Hydro- 
chloric acid neutralizes sodium hydroxide and forms sodium chlo- 
ride and water, (b) Nitric acid neutralizes potassium hydroxide 
and forms potassium nitrate and water, (c) Potassium chloride and 
silver nitrate form silver chloride and potassium nitrate, (d) Car- 
bon dioxide and water form carbonic acid (H 2 C0 3 ). (e) Barium 
oxide and water form barium hydroxide. (/) Potassium carbo- 
nate by hydrolysis forms acid potassium carbonate and potassium 
hydroxide. 

11. Write the ionic equations for (a) and (b) in Problem 10. 

12. Express the following reactions as ordinary chemical equa- 
tions and as ionic equations : (a) Sodium bromide and silver nitrate 
form silver bromide and sodium nitrate ; (b) ammonium iodide 
and silver sulphate form silver iodide and ammonium sulphate. 

13. Complete and balance : (a) Ag 2 S0 4 + KBr = K 2 S0 4 + ; 

(b) BaF 2 + H 2 S0 4 = HF + ; (c) NaBr + Cl 2 = Br + ; 

(d) H 2 S + = HI + ; (e) HI + = H 2 + I; (/) PBr, 

+ H 2 = H3PO3 + HBr. 

14. What weight of sodium hydroxide is needed to neutralize 
45 cc. of a sulphuric acid solution having the specific gravity 1.3 
and containing 32 per cent (by weight) of pure H 2 S0 4 ? 

15. In what proportions by weight do the following form a neu- 
tral solution : (a) Hydrobromic acid (HBr) and ammonium hy- 
droxide, (b) phosphoric acid (H3PO4) and calcium hydroxide, 

(c) sulphurous acid and sodium hydroxide? 

16. Calculate the weight of the salt formed in the following 
cases of neutralization : (a) Hydrochloric acid and 10 gm. of potas- 
sium hydroxide, (b) nitric acid and 37 gm. of barium hydroxide, 
(c) sodium hydroxide and 55 gm. of acetic acid. 

17. If 50 gm. of sodium nitrate (92 per cent pure) is made into 
nitric acid, what weight of sodium hydroxide (92 per cent pure) 
will neutralize the acid ? 

18. A ton of calcium acetate is needed. Calculate the required 
weight of lime (97 per cent CaO) and the volume of acetic acid 
solution having the specific gravity 1.035 and containing 25 per 
cent (by weight) of the pure acid. 



CHAPTER XI 
Energy and Chemical Change — Chemical Equilibrium 

Chemical change is always attended by the production or 
consumption of one or more of the different forms of energy, 
such as light, heat, and electricity. This means that a 
chemical change involves not only a rearrangement of matter, 
but also a transformation or redistribution of energy. Thus, 
when coal is burned, a new compound, called carbon dioxide, 
is formed, but heat is also liberated as a result of the 
chemical change. Sometimes we pay more attention to 
the redistributed matter than to the energy, but both 
are involved. In the present chapter we shall emphasize 
the relation of energy to chemical change. (See Chapter I.) 

The law of the conservation of energy should be recalled 
in this connection (see page 8). Energy cannot be created 
or destroyed by any known means ; we can only transform 
it. Hence, the chemical energy that is in elements and 
compounds appears as heat, light, or electricity when chemi- 
cal changes occur. And these forms of energy are in turn 
transformed into chemical energy and stored up in the 
chemical elements and compounds which are the result of 
chemical changes. 

Light and Chemical Change 

Light is often produced by chemical change. Sometimes 
the light is faint, as in the slow oxidation of yellow phos- 
phorus, which is luminous in moist air. This phenomenon 
is also exhibited by mixtures containing even a very little 

169 



170 INORGANIC CHEMISTRY 

phosphorus ; for example, the head of a phosphorus match. 
When the phosphorus and oxygen unite, part of the energy 
in each element is transformed into light and part is stored 
up in the phosphorus pentoxide. Usually the transforma- 
tion of chemical energy into light is more vivid ; that is, 
more chemical energy becomes light. Many chemical ex- 
periments are accompanied by intense light, especially those 
involving combination with oxygen or with chlorine. Thus, 
magnesium burns in oxygen with a dazzling light, and 
powdered antimony, as well as some other metals, bursts 
into a flame when dropped into chlorine gas. Combustion 
in general, especially of coal, oils, and gases containing 
compounds of carbon, is usually attended by light, and 
serves as an excellent illustration of the transformation of 
chemical energy into light. (See Chapter XVI.) 

Light is also often transformed into chemical energy. 
This transformation is typically illustrated by photographic 
processes. Paper, glass plates, and films coated with com- 
pounds of silver are blackened on exposure to the light; 
the compounds are changed chemically and light is trans- 
formed into chemical energy. Another transformation more 
or less familiar is that involved in the fading of colored 
fabrics, wall paper, and paintings. Light is also absolutely 
essential in the complex chemical changes involved in the 
growth of plants. The sunlight is stored up in the plants 
and is subsequently utilized by mankind when wood and 
coal are burned as fuel or vegetable matter is consumed as 
food. Certain chemical changes which proceed very slowly 
are hastened by light. Thus, hydrogen and chlorine gases 
when mixed in the dark do not unite perceptibly, but they 
combine slowly in diffused light and instantaneously in the 
direct sunlight. Similarly, a solution of chlorine (in water) 
evolves oxygen slowly in the dark but more rapidly in the 
light. 



HEAT 171 

Heat and Chemical Change 

Heat and chemical change are closely and definitely re- 
lated. A chemical change is almost invariably accompanied 
by the liberation or absorption of heat, usually the libera- 
tion. Vigorous and rapid reactions develop considerable 
heat and are also often attended by light, while the heat 
evolved by feeble or slow reactions is comparatively slight 
and sometimes can scarcely be detected. 

A familiar instance of the evolution of heat by chemical 
change is the slaking of lime. Lime is calcium oxide (CaO), 
and' when lime and water are mixed their chemical union 
produces sufficient heat to boil water and often to set fire 
to wood. Steam can be seen escaping from the boxes in 
which lime is being mixed with water in the preparation of 
plaster or mortar. Buildings in which lime is stored some- 
times take fire, if rain leaks in upon the lime. 'Ships loaded 
with lime are in constant danger of being set afire. An- 
other illustration is provided by the combustion of fuels 
such as coal, wood, oils, and gases. These are largely car- 
bon. The carbon in these substances unites with the oxygen 
of the air, and the chemical energy in both elements becomes 
heat to a great extent, some of course remaining in the 
products of combustion. Many chemical changes, already 
considered, are attended by the liberation of heat, the most 
conspicuous being the act of combination with oxygen. 
Thus, when hydrogen burns, the act of combination is 
strikingly manifested by the colorless, intensely hot flame. 
The interaction of metals and acids, as seen in the prepara- 
tion of hydrogen, develops heat. The chemical union of 
sulphur and metals is often accompanied by heat sufficient 
to cause incandescence. Reduction of metallic oxides often 
liberates much heat. For example, the heat attending the 
reduction of iron oxide by aluminium is so intense that the 
iron melts. (See Thermit.) 



172 INORGANIC CHEMISTRY 

Many chemical changes take place slowly at the ordinary 
temperature. Once started by heat, however, they proceed 
until the interacting substances are exhausted or the ex- 
ternal supply of heat is removed. There are many illus- 
trations of this hastening of chemical change by heat. 
Magnesium tarnishes very slowly in the air, but if a lighted 
match is applied to the metal, oxidation proceeds rapidly 
until the magnesium is entirely changed into magnesium 
oxide. Hydrogen and oxygen mix freely at the ordinary 
temperature without appreciable combination; if mixed in 
the proportion of two volumes to one and heated to 600° C, 
combination takes place in about one hour, while union in- 
stantly occurs when heat is applied in the form of a flame 
or an electric spark. Similarly, illuminating gas must be 
lighted before it will interact chemically with the oxygen of 
the air, but once raised to the proper temperature, the 
chemical change continues as long as the gas is supplied. 
Combustible substances, such as wood, many oils and gases, 
sulphur, and phosphorus, must be raised to a minimum tem- 
perature called the kindling temperature before the chemical 
changes attending combustion can proceed. The kindling 
temperature varies with the physical condition in the case of 
many solids, being lower, as a rule, when the substance is 
finely divided or presents a relatively large surface. Thus, 
shavings catch fire at a lower temperature than a stick or 
log of the same variety of wood. In many chemical changes 
heat must be constantly supplied. Thus, mercuric oxide 
decomposes into mercury and oxygen as long as it is suffi- 
ciently heated, but w T hen the flame is removed, the chemical 
change slackens and soon ceases altogether. The same is 
true of potassium chlorate. Indeed, heat is one of the 
most efficient aids to chemical change, and various sources 
of heat are indispensable in the laboratory as well as in 
chemical manufactories. 



HEAT 



173 



Sources of Heat. — Heat is so essential in all chemical 
operations that chemists have devised and used many ap- 
pliances for generating intense and continuous heat. The 
alchemists burned wood and charcoal in the furnaces which 
heated their crucibles. Priestley and Lavoisier employed a 
lens or burning glass in some of their experiments. Liebig 
and his contemporaries used a charcoal furnace in analyzing 
organic compounds. The greatest advance was made when 
Bunsen invented the burner which consumes gas and pro- 
duces a hot, smokeless flame. This burner is replaced by 
the blast lamp, oxyhydrogen blowpipe, or oxyacetylene blow- 
pipe when a high temperature is desired. In the arts and 
industries various kinds of furnaces are used ; e.g. rever- 
beratory, open hearth, and blast (see Figs. 64, 82, and 80). 
But all these sources of heat have been surpassed by the 
electric furnace. It is well 
known that an electric arc 
light produces intense heat. 
The high temperature of 
the arc, i.e. space between 
the glowing ends of the 
carbons, is unequaled by 
that of any other source 

r. , • r ' i i i tc ii Fig. 19. — Electric furnace — arc type. 

oi artificial heat. It the 

carbon rods are inclosed in an infusible box or vessel 
that prevents escape of heat, a temperature estimated to 
be about 3500° C. can be produced inside the receptacle. 
This apparatus is called an electric furnace. One type of 
the electric furnace is shown in Figure 19. When a cur- 
rent is passed through the carbon rods, the tremendous 
heat produced within the space is retained by the non- 
conducting walls and acts upon the substance in the cru- 
cible below the arc. The outside of the furnace remains 
cold enough to be touched by the hand, but the inside 




174 INORGANIC CHEMISTRY 

is almost twice as hot as the oxyhydrogen flame. There 
is no electrical action upon the chemicals. The intense 
heat alone produces the remarkable physical and chemical 
changes, and for this reason the process is often called an 
electrothermal process. Sand, lime, magnesium oxide, and 
other refractory substances melt and volatilize. The ele- 
ments carbon, silicon, and boron boil; and gold, copper, 
and platinum quickly melt and vaporize. Large masses of 
rare and uncommon elements are quickly reduced from 
their oxides and obtained in the pure state; e.g. chromium, 
manganese, tungsten, uranium, and molybdenum. Stable 
compounds of carbon, boron, and silicon are formed. These 
are the carbides, borides, and silicides. Some of the car- 
bides have an industrial use as well as scientific interest, 
especially calcium carbide and silicon carbide. 

Another type of electric furnace, known as the resistance 
type, is shown in Figure 20. It is essentially an insulated 
box-like structure of 
heat-resisting mate- 
rials. The wires con- ^^H ^^— — ^ ?|^VV^^^^^ 





veying the current are 
attached to perma- 

. ' , , Fig. 20. — Electric furnace — resistance type. 

nent outer ends and 

the carbon electrodes project into the furnace. Pieces of 
broken carbon make electrical connection between the elec- 
trodes and at the same time offer great resistance to the 
current. Hence intense heat is developed along the carbon 
core. Large electric furnaces constructed on this type are 
now in practical operation. And since electricity is obtained 
in many localities by operating dynamos by water, new in- 
dustries requiring intense and continuous heat have quickly 
sprung into existence and older ones have bfeen remodeled. 
Some of these plants are located at Niagara Falls, which fur- 
nishes enormous power at a relatively small expense. Several 



HEAT 175 

commercial substances, more or less familiar, are manu- 
factured in the electric furnace; e.g. calcium carbide, 
carborundum, alundum, phosphorus, carbon disulphide, 
graphite, and silicon. These are discussed in appropriate 
places. 

Measurement of Heat Energy. — Every substance possesses 
a certain amount of chemical energy, but there is no way of 
determining the total amount. We can, however, measure 
that part of the chemical energy which is transformed into 
heat when a substance or set of substances undergoes a 
chemical change. Thus, when hydrogen and oxygen com- 
bine chemically, the total amount of chemical energy in the 
two gases is divided, part being liberated as heat, and part 
being locked up as chemical energy in the water formed ; 
and the liberated heat may be taken as a measure of the 
chemical energy transformed during the chemical change. 
Heat measurements are made in a calorimeter. This appa- 
ratus consists essentially of two parts, a small vessel in which 
the substance is chemically changed and a larger one contain- 
ing water in which the small vessel is immersed. The heat 
involved in the chemical reaction changes the temperature of 
the water. The fuel (i.e. heat) value of food, coal, etc., is 
found by means of a calorimeter. Heat is measured in cal- 
ories; a calorie (cal.) is the quantity of heat necessary to 
raise 1 gm. of water 1° C. in temperature (15°-16° C. being 
the degree usually taken). For example, the heat liberated 
by the burning of 1 gm. of hydrogen is 34,200 cal. Atten- 
tion has already been called to the high temperature of the 
hydrogen flame. (See Chapter III.) 

Ordinary chemical equations do not express changes in 
energy. To represent heat changes, the number of calories 
of heat involved is placed after the equation, together with 
the proper sign, thus : — 



176 INORGANIC CHEMISTRY 

+ 68,400 cal. 



2H 


+ o = 


H 2 


Hydrogen 


Oxygen 


Water 


2 


16 


18 



This is called a thermal equation, and it means that 68,400 
cal. of heat are liberated when 2 gm. of hydrogen unite with 
16 gm. of oxygen to form 18 gm. of water. In some chemi- 
cal changes heat is absorbed. Thus, when carbon unites 
with sulphur to form carbon disulphide, heat is absorbed. 
The equation expressing this fact is — 

19,600 cal. 



c + s 2 = 


cs 2 


Carbon Sulphur 
12 64 


Carbon 

Disulphide 

76 



Heat evolved or absorbed in the formation of a mole of a 
compound is called its heat of formation. If heat is liber- 
ated in the formation of a compound, the heat is desig- 
nated positive (+) ; and the compound is termed exothermic. 
Heat of formation which is absorbed is designated negative 
( — ); and a compound having a negative heat of forma- 
tion is said to be endothermic. Exothermic compounds are 
relatively stable ; they can be decomposed by the addition of 
the same quantity of heat liberated by their formation. 
Thus, 68,400 cal. of heat or an equivalent quantity of energy 
must be added to 18 gm. of water to decompose it into 2 gm. 
of hydrogen and 16 gm. of oxygen. Such heat is called heat 
of decomposition. On the other hand, endothermic com- 
pounds are unstable and often explosive. They decompose 
easily with the liberation of heat. The heat evolved when 
organic substances are ultimately oxidized to carbon dioxide 
and water is called heat of combustion. Our knowledge of 
the calorific value (page 269) of fuels and foods is based 
largely on measurements of their heats of combustion. In 
the preceding chapter attention was called to the fact that 
the interaction of equivalent quantities of strong acids and 



ELECTRICITY 177 

bases liberates the same quantity of heat ; that is, the heat 
of neutralization liberated by the union of hydrogen and 
hydroxyl ions is 13,700 cal. under normal conditions. The 
thermal equation for the neutralization of sodium hydroxide 
by hydrochloric acid is — 

NaOH + HC1 = NaCl + H 2 + 13,700 cal. 

Applications and extensions of thermochemistry, as this 
branch of the science is called, will be made in subsequent 
chapters. 

Electricity and Chemical Change 

The Relation between Electricity and Chemical Action has 
always been a fascinating subject. Volta constructed his 
voltaic pile about 1800. This was one of the first (perhaps 
the first) sources of an electric current. In May, 1800, 
Nicholson and Carlisle decomposed water into hydrogen 
and oxygen by an electric current obtained from a thermo- 
pile. In the same year Cruikshank obtained lead and 
copper from solutions of their salts. And in 1807 Davy 
isolated two elements, sodium and potassium, by passing 
an electric current (obtained from a large battery) through 
fused sodium hydroxide and potassium hydroxide respec- 
tively. From that time until the present day the relation 
between electricity and chemical change has engaged the 
attention of chemists, and their labors have built up a 
branch of chemistry called electrochemistry, which has 
recently attained considerable commercial importance. 

Transformations of Electrical and Chemical Energy. — 

Electricity, like heat and light, is readily transformed into 
chemical energy, and vice versa. Certain chemical changes 
produce electricity, while electricity is consumed in others. 
A typical illustration of the transformation of chemical 



: 


Z G% 


^^^EE^z^E^ 


-~ 


n 


BBBBfifiJIS 


^^=^==^==^^ 



178 INORGANIC CHEMISTRY 

energy into electricity is furnished by the 
voltaic cell. The simplest form consists 
of two unlike metals, such as copper and 
zinc, connected by a wire and partly 
immersed in a vessel containing dilute 
sulphuric acid (Fig. 21). When the con- 
Fig. 21. — Simple vol- nected metals are put into the acid, the 

taic cell. Z is the z j nc s l wly disappears and hydrogen bub- 
zinc and C is the -, , ., ^ , , 

bles appear on the copper, further ex- 
amination shows that zinc sulphate is also 
formed. The chemical change is essentially the one already 
described under hydrogen, and can be represented thus : — 

Zn + H 2 S0 4 = 2H + ZnS0 4 

Zinc Sulphuric Hydrogen Zinc 

Acid Sulphate 

The wire becomes electrified and exhibits the effects of an 
electric current. For example, it becomes warm and makes 
a magnetic needle move. The source of the electric current 
is the chemical action between the acid and zinc. The 
copper is necessary, otherwise the product of the chemical 
action would be merely heat. Carbon is often used instead 
of copper, and sulphuric is replaced by other liquids, such 
as solutions of ammonium chloride or potassium hydroxide. 
The liquid chosen, however, must be one that will interact 
with zinc or its substitute. The zinc is ultimately trans- 
formed into zinc sulphate or some other chemical compound, 
and must be replenished; the solution must likewise be 
renewed. This necessity of recharging the cell is a striking 
proof of the intimate relation between chemical change and 
electricity. Several cells joined together form an electric 
battery. For many years the battery was the chief source 
of the electric current ; and it is now wddely used to generate 
the currents of moderate intensity utilized in ringing bells 



ELECTRICITY 179 

and operating electrical apparatus. Powerful currents, as 
a rule, are obtained from a dynamo. 

The transformation of electricity into chemical energy is 
illustrated by electrolysis. In Chapter IX considerable 
space was devoted to a discussion of the electrolysis of solu- 
tions; i.e. to the chemical changes which accompany the 
passage of an electric current through a solution. But 
electrolysis is not limited to solutions. Fused {i.e. melted) 
substances also undergo chemical changes when subjected 
to the action of a powerful electric current. As already 
stated, electrolysis is accomplished in an electrolytic ceil. 
This apparatus differs from a voltaic cell in one essential 
respect. The voltaic cell produces electricity, whereas the 
electrolytic cell consumes electricity. Otherwise both have 
three main parts, — the containing 
vessel, two electrodes, and the 
electrolytic solution or fused elec- 
trolyte. A simple electrolytic cell 
is shown in Figure 22. In such a 
cell the current from the battery 
or dynamo enters the cell through 
the positive ( + ) electrode or anode FlG - 2 " 2 - — Simple electrolytic 

i-i xl l xi 4.' cell. A is the anode, C is 

and leaves through the negative the cathodej and E ^ ^ 
( — ) electrode or cathode. Elec- trolytic solution. 
trodes may be of platinum, copper, 

zinc, mercury, or hardened carbon ; they may have any 
shape, — rod, wire, sheet, plate, box, crucible; and they 
also may be solid, liquid, or powder, as well as fixed or 
movable. The electrodes are connected by wdres with the 
source of the electric current, and serve as "doors" — 
to quote Faraday — for the current to flow into and out 
of electrolytic solution or the electrolyte. We speak of a 
"current" of electricity and of electricity as "flowing," 
although we do not know the fundamental nature of elec- 





t 


' : +A C- : - 


p 


—7- _, 


-^^^=-=r^=^z .=_—-- 


=i 






"z lV 


. llSglljliifflgii t 


H 




E^^^H-H^^^E^Z^^^^^E^E^E^^^^^ 







180 INORGANIC CHEMISTRY 

tricity nor do we mean really that it flows, like a river, only 
in one direction. The anode is the electrode that is often 
consumed or worn away, either mechanically or chemically, 
but solids, especially metals, are often deposited upon the 
cathode. The containing vessel likewise may have any 
desired shape, and is usually made of some material which 
will resist the corrosive action of chemicals; e.g. porcelain, 
slate, or soapstone. Metallic vessels lined with various sub- 
stances selected to withstand intense heat are also used, 
since fused electrolytes are often subjected to electrolysis. 

The phenomena exhibited during the electrolysis of solu- 
tions have been described. (See Chapter IX.) Certain 
features of the operation, however, should be recalled at 
this point, since they apply to both fused and dissolved 
electrolytes. Employing the interpretation offered by the 
theory of electrolytic dissociation, the ions of the electrolyte 
begin to migrate as soon as the electric current enters the 
cell; the electro-negative anions move toward the electro- 
positive anode, and the electro-positive cations move toward 
the electro-negative cathode. Reaching their respective 
electrodes, the ions give up their charges and become atoms, 
atomic groups, or molecules. The discharged particles may 
escape as gases, dissolve in the liquid, or attach themselves 
to the electrodes; very often secondary chemical changes 
occur which complicate the process and sometimes cause 
serious difficulties in industrial applications of electrolysis. 

A third illustration of the transformation of electrical 
energy into chemical energy, and vice versa, is furnished by 
the storage cell. This cell consists essentially of two grids 
of lead which are filled with a mixture of sulphuric acid and 
lead oxide (PbO) and then immersed in a vessel containing 
dilute sulphuric acid. The mixture in the grids soon be- 
comes lead sulphate (PbS0 4 ). When an electric current is 
passed through the cell from one plate to the other, the 



ELECTRICITY 181 

hydrogen ions formed by the ionization of the sulphuric 
acid migrate to the cathode, where they become atomic 
hydrogen and by their chemical interaction with the lead 
sulphate form sulphuric acid and metallic lead — the latter 
remaining attached to the grid (cathode); the S0 4 -ions 
migrate to the anode, where they become ordinary uncharged 
atomic groups and by their interaction with the lead sul- 
phate form sulphuric acid and lead dioxide (Pb0 2 ) — the 
latter remaining attached to the grid (anode). These 
changes continue until nearly all the lead sulphate has 
been altered as just described. If, after being charged, the 
grids are connected by a wire, a current of electricity will 
be obtained in a direction opposite to that used in charging 
the cell, and the chemical changes will take place in the 
reverse direction. The cell gradually reverts to its original 
condition, and must be recharged if a current is again 
desired. In this cell, therefore, there is a complete circuit 
of transformation, — electricity — > chemical energy — > elec- 
tricity. 

Industrial Applications of Electrolysis. — The earliest in- 
dustrial application of electrolysis was in electrotyping and 
electroplating. These operations consist in depositing a thin 
film of metal upon a surface. They are fundamentally the 
same, though copper is the only metal used for producing 
electrotypes. Electrotypes are exact reproductions of the 
original object. The process of electrotyping is substantially 
as follows : The page of type, for example, is first repro- 
duced in wax. This exact impression is next covered with 
powdered graphite to make it conduct electricity. The 
coated mold is then suspended as the cathode in an acid 
solution of copper sulphate; the anode is a plate or bar of 
copper. When the current is passed through the system, 
electrolysis occurs; copper is dissolved from the anode and 



182 INORGANIC CHEMISTRY 

deposited on the mold in a film of any desired thickness. 
The exact copper copy is stripped from the mold, backed 
with metal, and used instead of the type itself. By this 
process, exact copies of expensive wood engravings can be 
cheaply reproduced and type can be saved from the wear 
and tear of printing. Most books, magazines, and news- 
papers are now printed from electrotypes. The process of 
electroplating differs from electrotyping in only one essen- 
tial; viz. in electroplating the deposited film is not removed 
from the object. The object to be plated is carefully cleaned 
and made the cathode; the anode is a bar or plate of the 
metal to be deposited. When the current passes through 
the system, the metal is firmly deposited on the object. 
The electrolysis would take place, of course, if any anode 
were present; but anodes of the metal to be deposited are 
usually used to prevent the solution or "bath" from weaken- 
ing. They accomplish the purpose by replenishing the solu- 
' tion with metal as fast as it is removed and deposited upon 
the cathode. Silver, nickel, and gold are the usual metals 
used in electroplating. (See these metals.) 

Electroplating and electrotyping have been done since 
about 1840. It is only within the last ten or fifteen years, 
however, that the electric current has been profitably ap- 
plied in many industries. But during this time the develop- 
ment of electrochemistry has been very marked. The largest 
of these industries is the refining of copper. The process is 
similar to that described under electrotyping. Other metals, 
such as gold, silver, and lead, are extracted from their ores 
and purified by electricity, though the older processes are 
still used. The aluminium, magnesium, and sodium of 
commerce are now manufactured by passing an electric 
current through their fused compounds. Chlorine, potas- 
sium chlorate, potassium hydroxide, and sodium hydroxide 
are some of the other important industrial products of 



ELECTRICITY 183 

electrolysis. These electrochemical processes and products 
will be fully discussed in the appropriate places. 

Measurement of Electrical Energy. — Faraday was the 
first scientist to make a thorough study of electrolysis. He 
found that a given current of electricity liberated different 
but definite amounts of the chemical elements. Thus, 
the current which liberates 1 gm. of hydrogen also liber- 
ates 8 gm. of oxygen, 35.46 gm. of chlorine, 107.88 gm. of 
silver, 31.78 gm. of copper, and so on. These numbers are 
identical with the chemical equivalents of these elements. 
(Compare Equivalents, Chapter XIV.) Faraday called them 
electrochemical equivalents, to emphasize their chemical 
and electrical relationship. But the term electrochemical 
equivalent now means the w T eight of an element deposited 
or liberated by a specified current in a certain time (1 ampere 
in 1 second). For example, the electrochemical equivalent 
of hydrogen is .000010441 gm., of oxygen is .00008287 
and sometimes .00016574, of copper is .0003294 and some- 
times 0.0006588, of silver is .001118. This general relation 
is often stated as Faraday's law, thus : — 

When the same current of electricity is passed through solu- 
tions of different electrolytes, the ratio of the quantities of 
liberated products is the same as that of their chemical equiva- 
lents. 

Faraday also showed that — 

The amount of decomposition — the chemical work, we 
might say — is proportional to the total amount of electricity 
used. It makes no difference whether the current is strong 
or weak, nor whether the time of its flow is long or short. 
A certain quantity of electricity will do so much chemical 
work — no more and no . less. Thus, a given quantity of 
electricity passed through copper sulphate solution always 
deposits the same weight of copper at the cathode. 



184 INORGANIC CHEMISTRY 

These two principles of Faraday are at the foundation 
of all electrochemical industries. Their importance can 
hardly be overestimated. 

Other Conditions affecting Chemical Action 

Substances vary in their tendency to undergo chemical 
change. Some, like oxygen, are active elements ; i.e. they 
unite directly with other substances, liberate energy rapidly, 
and form comparatively stable compounds. Others, like 
nitrogen, are inert, and form somewhat unstable compounds 
by indirect processes. Not only does the chemical activity 
of substances differ, but the same substance is not equally 
active toward all others. Thus, iron unites readily with 
sulphur and bromine, but not with copper or mercury. 
Nevertheless, certain substances which are inert under some 
conditions may become active under special conditions. 
That is, chemical action depends not only upon the specific 
attraction of interacting substances, i.e. their chemical 
affinity as it is sometimes called, but also upon the special 
conditions under which the reaction occurs. Chemical 
changes which take place very slowly, as we have already 
seen, sometimes proceed with astonishing rapidity under the 
influence of heat, light, or electricity ; that is, the velocity of 
the reaction is increased. In many chemical changes, how- 
ever, the velocity of the reaction is not influenced solely by 
one of the familiar forms of energy, but also by other factors, 
e.g. solution, concentration, catalysis, and equilibrium. By 
the velocity of a reaction we mean the amount of substance 
that reacts in a given time. 

The Effect of Solution on chemical change was discussed 
in Chapter IX. Many chemical changes are carried out in 
aqueous solutions because dissolved substances are in a con- 
dition especially favorable for interaction. Therefore " wet 



CATALYSIS 



185 



processes/ ' as the}^ are sometimes called, are exceedingly 
important in chemical analysis. The velocity with which a 
reaction proceeds is often greatly influenced by concentration ; 
that is, by the quantity of substance in a given volume. If 
the concentration is increased, the reaction takes place more 
rapidly; that is, in general, the greater the concentration, 
the greater the reaction velocity. A concrete illustration 
will make this point clearer. An acid solution of potassium 
iodide and starch turns blue when an oxidizing agent is 
added, owing to the interaction of the starch and the liber- 
ated iodine. If the potassium iodide solution and oxidizing 
solution are added to the same volume of water, the time 
required to produce a standard blue color will be a relative 
measure of the reaction velocity. The following table shows 
the result of an experiment : — 

Measurement of Reaction Velocity 



Solution 


Oxidizing Solution 


Potassium Iodide 
Solution (oc.) 


Time in Minutes 


I 

II 
III 
IV 


5 
10 

5 
10 


5 

5 

10 

10 


4 

2 

2 
1 



From the table it is evident that as the concentration of 
solution IV is four times that of solution I, the time for IV 
is one fourth that of I ; solutions II and III have the same 
concentration and their times are the same. 



Catalysis. — The velocit} 7 of some reactions can be altered 
by the presence of certain substances that apparently do 
not participate in the chemical change and can be recovered 
unchanged after the action has ceased. Such substances 
are called catalytic agents or catalyzers, and their effect 



186 INORGANIC CHEMISTRY 

upon chemical action is called catalysis or catalytic action. 
For instance, potassium chlorate yields oxygen slowly when 
heated to about 350° C, but if powdered manganese dioxide 
is added, the gas is evolved rapidly. Furthermore, the 
manganese dioxide can be recovered after the experiment 
merely by dissolving out the residual potassium chloride. 

Similarly, hydrogen dioxide (H 2 2 ) decomposes very 
slowly in the air, but if manganese dioxide is added, the 
decomposition proceeds rapidly and bubbles of oxygen gas 
can be seen rising through the liquid. Likewise, a mixture of 
hydrogen and oxygen at the ordinary temperature reveals no 
tendency toward chemical action, but if a little powdered 
platinum is added, the gases combine with almost explosive 
violence. Again, sulphur dioxide gas (S0 2 ) and oxygen 
united very slowly, but when a purified and properly cooled 
mixture of these gases is passed over finely divided platinum, 
the gases unite rapidly and form sulphur trioxide (S0 3 ). 
This rapid transformation is utilized in one process of manu- 
facturing sulphuric acid. Water vapor in traces is regarded 
by some authorities as a catalytic agent. Thus, many gases, 
especially hydrogen chloride (HC1) and ammonia (NH 3 ), do 
not unite when perfectly dry, but if a trace of water is added, 
the reaction proceeds as ordinarily observed ; i.e. white 
fumes of ammonium chloride (NH 4 C1) are formed. In many 
technical processes reactions are hastened by catalyzers, 
often in relatively small quantities, e.g. mercury, nickel, 
chlorides of copper, zinc, and aluminium, dilute acids, and 
enzymes. Sometimes a catalyzer retards a reaction ; such 
substances are called negative catalyzers. Catalysis is an 
important phenomenon. 

Reversible Reactions and Chemical Equilibrium. — Hitherto 
we have described chemical reactions as if they proceeded 
to completion. Indeed, they do apparently, for in many 



CHEMICAL EQUILIBRIUM 187 

experiments conditions are chosen which necessitate action 
until one or more of the interacting substances is exhausted. 
As a matter of fact many reactions are reversible; that is, 
they proceed in one direction under one set of conditions 
and in the opposite direction under another set of condi- 
tions. For example, in the manufacture of oxygen from 
barium oxide a reversible reaction occurs. When barium 
oxide is heated in the air to about 700° C, the chemical 
change is represented thus : — 



BaO + = 


Ba0 2 


Barium Oxygen 
Oxide 


Barium 
Dioxide 



If the air supply is cut off and the pressure (in the retorts) 
is reduced, the chemical change is reversed and may be 
represented thus : — 



Ba0 2 


= O + BaO 


Barium 
Dioxide 


Oxygen Barium 
Oxide 



Another illustration is provided by Lavoisier's famous ex- 
periment described under Hydrogen (see Chapter III). He 
passed steam over red-hot iron and obtained hydrogen and 
iron oxide; the equation for the chemical change is — 

4H 2 + 3Fe = 8H + Fe 3 4 

Water Iron Hydrogen Iron 

Oxide 

When hydrogen is passed over hot iron oxide, the chemical 
change is reversed, thus : — 

8H + Fe 3 4 = 3Fe + 4H 2 

Hydrogen Iron Iron Water 

Oxide 

We might conclude that in reversible reactions the chemi- 
cal change can proceed in either direction to completion. As 
a matter of fact the actual or effective chemical change de- 
pends upon the conditions. Thus, in the interaction of 



188 INORGANIC CHEMISTRY 

steam and iron, if the hydrogen is continuously swept out 
of the tube by the steam, no reduction of the iron oxide 
occurs ; or if in the interaction of hydrogen and iron oxide 
the steam is continuously removed, no oxidation of iron will 
occur. But when the experiment is performed in a closed 
tube, the result is an equilibrium between the two oppos- 
ing reactions. Not only does the interaction of the steam 
and iron produce hydrogen and iron oxide, but conversely, 
the interaction of the hydrogen and iron oxide forms steam 
and iron. The reactions proceed simultaneously in the same 
tube until chemical equilibrium is reached ; that is, not to 
completion in either direction, but to such a point in both 
directions that there is no further alteration in the weights 
of the substances actually participating in the change. Let 
us illustrate more definitely. If iron and steam are heated 
in a closed tube, the velocity of the forward reaction (i.e. the 
transformation of steam and iron into iron oxide and hydro- 
gen) gradually diminishes until it apparently stops, although 
some material is still available for chemical action ; con- 
versely, the velocity of the reverse reaction (i.e. the reforma- 
tion of steam and iron from iron oxide and hydrogen) gradu- 
ally increases until it likewise apparently stops, although 
additional material is available. Continued heating pro- 
duces no change in the relative weights of the four sub- 
stances in the tube ; that is, equilibrium has been reached. 
This means that the velocities of the opposing reactions are 
equal. In other words, there is no further accumulation of 
any of the reacting substances, because the opposing re- 
actions are proceeding at the same rate ; both the forward 
and reverse reactions are still taking place, but one undoes 
the work of the other, so to speak. Chemical equilibrium, 
then, is that state reached in a reversible reaction when the 
further accumulation of any of the reacting substances is pre- 
vented by the equal velocities of the two opposing reactions, 



CHEMICAL EQUILIBRIUM 189 

In equations representing reversible reactions oppositely 
directed arrows (read " equals reversibly ") are used, thus : — 

H 2 + I 2 ^± 2 HI. 

The velocity of a reaction is affected by several factors, 
especially the concentration of the reacting substances. 
The law covering the relation between the concentration 
and the velocity of a reaction is often called the law of mass 
action, and may be stated thus : 

At a constant temperature the velocity of a reaction is pro- 
portional to the molecular concentration of each reacting 
substance. 

By concentration of the reacting substance is meant that 
relative portion of each substance actually available for 
chemical action at a given time. By molecular concentra- 
tion is meant the number of moles (i.e. gram-molecular 
weights) in a liter. The law finds its best application in the 
case of gases and solutions because these systems of sub- 
stances, being homogeneous mixtures, are in a favorable 
condition for chemical action ; moreover, the concentration 
of each ingredient of a homogeneous mixture can readily be 
expressed. In applying the law to heterogeneous mixtures, 
e.g. steam-iron-iron oxide-hydrogen, it is customary to con- 
sider only the gases because the concentration of the par- 
ticipating portion of each solid — the active mass, as it is 
sometimes called — is practically constant. 

Let us apply the law of mass action to a gaseous reaction, 
viz. the combining of hydrogen and iodine to form hydriodic 
acid. The equation correctly written for our present pur- 
pose is : — 

H 2 + I 2 = 2 HI. 

Let (a) and (b) represent the molecular concentration of the 
hydrogen and iodine respectively. The velocity of the re- 



190 INORGANIC CHEMISTRY 

action is proportional to each molecular concentration and 
therefore to their product. The velocity also depends on 
such factors as the specific affinities of the reacting substances, 
temperature, and catalytic agents. At a given temperature 
the effect of this complex set of factors is constant, and the 
combined effect is often called the affinity constant. The 
affinity constant has a different value for each reaction and 
is represented by k y k h etc. Hence the equation for the 
velocity of the reaction between hydrogen and iodine be- 
comes : — 

(1) velocity = (a) X (6) X k. 

Let us next consider the reverse reaction, viz. the decom- 
position of hydriodic acid into hydrogen and iodine, which 
is properly represented by the equation : — 

2 HI = H 2 + I 2 . 

In expressing the velocity of the decomposition of hydriodic 
acid, two facts must be taken into account : (1) two 
molecules of hydriodic acid (2 HI) are formed from a single 
molecule each of hydrogen (H 2 ) and iodine (I 2 ), and (2) the 
affinity constant has a different value from that used in the 
forward reaction. Consequently, if the molecular concen- 
tration cf hydriodic acid is represented by (c), then the 
velocity of the reaction is proportional to (c) 2 and to k x ; 
and the equation becomes : — 

(2) velocity = (c) 2 X ki. 

Since at equilibrium the velocities of the forward and reverse 
reactions are equal, then : — 

(a) X (6) X k = (c) 2 X ki. 
Therefore 

(a) X (6) = fci = R 
(c) 2 k 



CHEMICAL EQUILIBRIUM 191 

This new constant K is called the equilibrium constant. 
Since it is the ratio of the two affinity constants, its numeri- 
cal value is constant at a constant temperature. That is, 
whatever the original concentrations of the reacting sub- 
stances, reactions will occur such that at equilibrium the 
concentrations wdll give the same value for the constant K. 

Chemical equilibrium is a sensitive relation between 
weights of substances and it is easily influenced by a change 
in the conditions. In other words, a change in temperature 
or concentration favors one of the opposing reactions and 
thereby produces another state of equilibrium which corre- 
sponds to the new conditions. The change in equilibrium 
caused by a change in conditions is called displacement of 
equilibrium. The two main factors that cause displacement 
of equilibrium are temperature and concentration. A rise 
in temperature increases the velocity of a reaction ; the 
velocity in many cases is doubled by a rise of ten degrees. 
As a rule, however, the velocities of the opposing reactions 
in a reversible reaction are affected quite differently by a 
change in temperature. Hence equilibrium is usually dis- 
placed by a change in temperature. 

The direction in which equilibrium is displaced by a 
change in temperature depends on which reaction — the 
forward or the reverse — absorbs heat. The reaction that 
absorbs heat is favored by rise of temperature, and vice 
versa; thus in the case of the reversible reaction, 

4 H 2 + 3 Fe ^t 8 H + Fe 3 4 , 

the reverse reaction (indicated by the lower arrow) is favored 
because it absorbs heat. The relation of temperature and 
displacement of equilibrium is a special case of a broad 
generalization known as Le Chatelier's Law (or theorem), 
which may be stated thus : — 

When a system in equilibrium is subjected to a change of 



192 INORGANIC CHEMISTRY 

conditions (such as temperature and pressure) the system alters 
in the way that neutralizes, or tends to neutralize, the effect of 
the change. 

The influence of a change in concentration on equilibrium 
is apparent from a consideration of a concentration equation, 
e.g. the equation expressing equilibrium in the reversible 
reaction involving hydrogen, iodine, and hydriodic acid : — 

(a) X (6) 



(c)> 



= K. 



Suppose the concentration represented by (a) is changed, 
then the numerical value of the numerator will be changed 
and reactions must take place until the concentrations are 
such that the new state of equilibrium will give the same 
value of the constant K. In other words, equilibrium is 
displaced by changing concentration. Change in concen- 
tration may be brought about in several ways. One way 
is the removal of one product of the reaction. (1) If the 
iodine from the decomposition of hydriodic acid is partly 
concentrated by cooling one end of the tube, this iodine no 
longer participates in the reaction, and the concentration 
of the active iodine will be diminished. That is, equilibrium 
is displaced and the decomposition of hydriodic acid must go 
much further before a new state of equilibrium is estab- 
lished. (2) In the experiment with iron and steam, if the tube 
is opened and a current of steam is introduced, the hydrogen 
which is formed by the forward reaction will be removed and 
the reaction will continue until the iron is oxidized. That 
is, the removal of the hydrogen displaces the equilibrium 
and thereby permits the forward reaction to proceed to com- 
pletion. Similarly, if hydrogen is introduced into the tube, 
the steam will be removed and the reduction of the iron by 
the hydrogen proceeds to completion. In other words, the 
entire removal of one product of the reaction displaces the 






CHEMICAL EQUILIBRIUM 193 

equilibrium and allows the reaction to proceed to completion. 
(3) Removal of one product of a reaction is readily accom- 
plished by establishing certain reactions in solutions. If 
one of the products is insoluble, it is removed from the 
sphere of action. When silver nitrate and hydrochloric 
acid are dissolved in water, the solution contains at first 
the ions Ag + , N0 3 ~, H + , and CI", and undissociated mole- 
cules of silver nitrate (AgN0 3 ) and hydrochloric acid (HC1). 
But the ions Ag + and CI" at once form molecules of insoluble 
silver chloride (AgCl), thereby displacing the equilibrium 
between the ions and their undissociated molecules. The 
removal of these ions permits the dissociation of more and 
more silver nitrate and hydrochloric acid molecules, and the 
precipitation of silver chloride continues until the silver or 
chlorine ions are practically exhausted. (4) A gaseous or 
readily volatilized product of a reaction can be removed by 
raising the temperature ; two typical examples are the prep- 
aration of hydrogen chloride and nitric acid (see pages 
204, 217). 

Another way to change concentration and thereby dis- 
place equilibrium is actually to change the quantity of 
(1) the solvent or of (2) an ionic substance in a given solu- 
tion. Consider a sodium chloride solution. The solution 
contains molecules (NaCl) and ions (Na + and CI"), and 
the molecules are in equilibrium with the ions, thus : — 

NaCl ^ Na + + CI". 

The degree of dissociation of the molecules depends on the 
concentration of the solution (and the temperature). The 
molecular concentration of the sodium chloride may be 
represented by C and the ionic concentration by Ci (for 
Na + ) and C 2 (for CI") . Then we may write : — 

Cl X C-2 _ T7- 

c ~ 



194 INORGANIC CHEMISTRY 

(1) Now if water is added to the solution, some of the mole- 
cules will dissociate into ions, and the molecular and ionic 
concentrations will diminish as dilution proceeds, though the 
ionic concentration will diminish less rapidly than the molec- 
ular. On the other hand, if w T ater is removed from the 
solution by evaporation, the concentrations will increase — 
the molecular concentration in this case increasing more 
rapidly than the ionic. Equilibrium is displaced in both 
operations — diluting and concentrating — by the changes 
in concentration, because the reactions take place until the 
concentration fraction has the same value for K. 1 (2) Sup- 
pose hydrogen chloride (or concentrated hydrochloric acid) 
is added to a solution of sodium chloride. The latter solu- 
tion originally contained sodium chloride molecules (NaCl), 
sodium ions (Na + ), and chloride ions (Cl + ). The hydrogen 
chloride (or hydrochloric acid) provides hydrogen chloride 
molecules (HC1), hydrogen ions (H + ), and chloride ions 
(CI"). Some of the sodium ions must unite with chloride 
ions to form sodium chloride molecules in order to maintain 
equilibrium in the solution. Now if the molecular concen- 
tration of the sodium chloride was very large in the first 
solution, the newly formed sodium chloride molecules will 
be in excess of the amount that can dissolve in the water, 
and hence some sodium chloride must be precipitated. 

A method of displacing equilibrium similar to (2) above 
is quite effective in the case of substances which are not 
very soluble in water. Such substances readily form satu- 
rated solutions of rather small molecular concentration. 
Hence the equilibrium between the molecules and ions is 
readily displaced. Thus, although barium sulphate is very 
slightly soluble in water, it dissolves to some extent ; and 
the equilibrium equation may be written : — 

1 The value of K is not constant in the case of strong electrolytes because 
certain factors, not yet understood, cannot be incorporated in the equation. 



CHEMICAL EQUILIBRIUM 195 

(Ba ++ ) X (S0 4 — ) K 
(BaS0 4 ) 

Now when barium chloride solution is added to sodium sul- 
phate solution, barium sulphate is precipitated because a 
saturated solution of barium sulphate is almost instantly 
formed and all additional barium sulphate must be precipi- 
tated. This precipitation may be interpreted in another 
way. In a saturated solution at a given temperature the 
concentration of the solute (in this case barium sulphate) is 
constant (Ki) . Hence the equation may be written : — 

(Ba ++ ) X (SO4-) = K X Kx = K'. 

In other words, the product of the ionic concentrations in a 
saturated solution is constant (K'). This new constant 
(K 7 ) is called the solubility product of barium sulphate. 
When the product of the ionic concentrations in any solu- 
tion exceeds the numerical value of the solubility product, 
some of the ions unite to form a precipitate. As a rule, 
when solutions of two electrolytes are mixed, a double de- 
composition involving precipitation takes place if the product 
of the concentrations of any two ions exceeds the solubility 
product of the salt formed by their combination. 

Problems 

1. Calculate the affinity constant of an acetic acid solution hav- 
ing a total molar concentration of .1 and a molar concentration of 
each ion of .0013. Ans. .0000171. 

2. The molar solubility of potassium chlorate at 18° C. is .52 
and the per cent of ionization is 70 (i.e. .70). Calculate the solu- 
bility product. Ans. .13. 

3. If the molar solubility of barium sulphate is .00001 and the 
ionization is .98, what is the solubility product? 

4. Calculate the solubility product of calcium hydroxide if its 
molar solubility is .02 and its ionization is .88. 



CHAPTER XII 



Chlorine and Hydrochloric Acid 

Chlorine is an important gaseous element, and its com- 
pounds are useful, especially hydrochloric acid, sodium 
chloride, and bleaching powder. 

Occurrence. — Free chlorine is never found in nature, 
but its compounds are widely distributed, the most abun- 
dant being sodium chloride. Many 
compounds of chlorine with potas- 
sium, magnesium, and calcium are 
found in the deposits at Stassfurt in 
Germany. (See Potassium.) About 
2 per cent of the total amount of 
matter in the ocean is chlorine, and 
the salts found in sea water con- 
tain about 55 per cent of chlorine. 
Silver chloride — " horn " silver — 
is mined as an ore in the United 
States and Mexico. 

Preparation. — Chlorine is pre- 
pared in the laboratory by heating 

FiG.23.-Apparatusforpre- ft mixture of mangane se dioxide 
paring chlorine. The gas . . 

is generated in A, passes and hydrochloric acid (Fig. 23). 
through CE to the bottom This method was used by Scheele, 

of the bottle G and dis- who disc0vered the gas in 1774 . 
places the air, which escapes 

through a hole in the cover The equation" for the preparation 
F- of chlorine is — 

196 




CHLORINE AND HYDROCHLORIC ACID 197 

Mn0 2 + 4HC1 = 2 CI + MnCl 2 + 2H 2 

Manganese Hydrochloric Chlorine Manganese Water 

Dioxide Acid Dichloride 

This is an oxidizing process, since the hydrogen of the 
hydrochloric acid is oxidized to water, thereby liberating 
a part of the chlorine of the acid as free chlorine gas. 
Other oxidizing substances besides manganese dioxide may 
be used, such as potassium chlorate (KC10 3 ), potassium 
dichromate (K 2 Cr 2 7 ), red lead (Pb 3 4 ), or potassium per- 
manganate (KMn0 4 ). 

Sometimes chlorine is prepared in the laboratory by heat- 
ing a mixture of manganese dioxide, sodium chloride, and 
sulphuric acid. This method is substantially the same as 
the other, since a mixture of sulphuric acid and sodium 
chloride yields hydrochloric acid. The equation for this 
method of preparing chlorine is — 

2 H 2 S0 4 + 2 NaCl + Mn0 2 = 2 CI + Na 2 S0 4 +MnS0 4 +2 H 2 

Sulphuric Sodium Manganese Chlorine Sodium Manganese Water 
Acid Chloride Dioxide Sulphate Sulphate 

Chlorine is manufactured by several, processes, one of which 
(Deacon process) involves the same chemical change (oxidation) as 
the laboratory method. In the Deacon process hydrochloric acid is 
oxidized by oxygen obtained from the atmosphere. A mixture of 
hydrochloric acid gas and air is heated to 375° C. and passed through 
iron tubes containing balls of clay or pieces of brick previously satu- 
rated with copper chloride (CUCI2). The essential chemical change 
is the oxidation of the hydrogen of the hydrochloric acid, and it may 
be represented by the equation — 

2HC1 + = 2 CI + H 2 

Hydrochloric Oxygen Chlorine Water 

Acid 

The copper chloride acts as a catalytic agent in this process, and may 
be replaced by other chlorides. 

In the Weldon process, an impure native manganese dioxide, known 
as pyrolusite, is treated with hydrochloric acid in large earthenware 
retorts or stone tanks heated by hot water or steam. When no more 
chlorine is liberated, the residue is mainly manganese dichloride. 



198 INORGANIC CHEMISTRY 

This "still-liquor" was formerly thrown away, but by the Weldon 
process it is changed into manganese compounds, which are used to 
prepare more chlorine. (See Manganese Dioxide.) 

Chlorine is also prepared on a large scale by an electrolytic process. 
A solution of sodium chloride is electrolyzed in properly constructed 
electrolytic cells, and the chlorine which is liberated at the anode is 
conducted off through pipes. Sodium hydroxide is produced at the 
same time, and the process will be described under this compound. 

Properties. — Chlorine is a greenish yellow gas. Its color 
suggested the name chlorine (from the Greek word chloros, 
meaning greenish yellow), which was given to it about 1810 
by Davy, who spent several years in studying this gas and 
its compounds. It has a disagreeable, suffocating odor, 
which is very penetrating. If breathed, it irritates the 
sensitive lining of the nose and throat, and a large quantity 
would doubtless cause death. It is heavier than the other 
elementary gases, and is about 2.5 times heavier than air. 
Hence it is easily collected by conducting it to the bottom 
of a bottle and allowing it to displace the air ; the term 
downward displacement is sometimes applied to this method 
of collecting a gas. A liter of dry chlorine at 0° C. and 760 
mm. weighs 3.22 gm. 

Chlorine is moderately soluble in water, about three liters 
of the gas dissolving in one. liter of water under ordinary 
conditions. The solution is yellowish, smells strongly of 
chlorine, and is frequently used in the laboratory as a 
substitute for the gas. Chlorine water, as the solution is 
called, is unstable even under ordinary conditions, and 
must be kept in the dark. If the solution is placed in the 
sunlight, oxygen is soon liberated and hydrochloric acid is 
formed. Intermediate changes doubtless occur; but the 
simplest equation for the essential change is — ■ 

H 2 + 2 CI = 2 HC1 + O 

Water Chlorine Hydrochloric Oxygen 

Acid 






CHLORINE AND HYDROCHLORIC ACID 199 

Chlorine is much less soluble in a solution of sodium chloride, 
over which it is sometimes collected. It attacks mercury, 
and cannot be collected over this liquid. 

Chlorine hydrate is formed by cooling concentrated chlorine water, 
or by passing chlorine into ice water. It is a yellowish, crystalline 
solid, and in the air it decomposes quickly into chlorine and chlorine 
water. Its composition corresponds to the formula CI2 . 8 H 2 0. 

Liquid Chlorine was first prepared by Faraday in 1823. A little 
chlorine hydrate was inclosed in one arm of a bent tube, which was 
then sealed. By gently heating the tube, the chlorine hydrate was 
decomposed into chlorine and water, but the chlorine, being unable 
to escape, was condensed to a liquid by its own pressure inside the 
tube. The liquefaction is more easily accomplished, if one end is 
kept cold during the experiment. This method has .been replaced 
by a simpler one; viz. subjecting the gas to a high pressure and 
moderately low temperature. The critical temperature is +141° C. 
and the critical pressure is 84 atmospheres. It is an oily, yellow 
liquid and is now a common commercial substance. Solid chlorine 
is a yellow crystalline mass. 

Chlorine is an active chemical element. It unites directly 
with most of the elements, the only conspicuous exceptions 
being oxygen, carbon, and nitrogen. Chlorine does not 
burn in the air, because it does not form an oxide directly. 
Many elements unite vigorously with chlorine. Thus, the 
metals antimony and arsenic, when sprinkled into chlorine, 
suddenly burst into flame, while the non-metal phosphorus 
melts at first and finally burns with a feeble flame. If 
sodium, iron powder, copper wire, or other metals are heated 
and then put into chlorine, they burn; the sodium and iron 
produce a dazzling light, and the copper glows and emits 
dense fumes of whitish smoke. The compound formed 
in each case is a chloride, i.e. a compound of chlorine and one 
other element. (See Chlorides, below.) Chlorine combines 
readily with hydrogen. Hence, a jet of burning hydrogen 
when lowered into chlorine continues to burn, forming 



200 INORGANIC CHEMISTRY 

hydrochloric acid gas, which appears as a white cloud. The 
simplest equation for this change is — 

H + CI = HC1 

Hydrogen Chlorine Hydrochloric 
Acid 

A mixture of hydrogen and chlorine explodes violently when 
exposed to the sunlight. Many compounds of hydrogen 
are decomposed by chlorine. Thus, compounds of hydrogen 
and carbon, such as those found in illuminating gas, paraffin 
wax, and wood, burn in chlorine with a smoky flame ; since 
chlorine does not combine directly with carbon, the flame 
contains multitudes of very fine particles of solid carbon. 
A piece of glowing charcoal is extinguished by chlorine. 
If cotton is saturated with warm turpentine (Ci H 16 ) and 
put into a bottle of chlorine, a flame accompanied by a dense 
cloud of black smoke bursts from the bottle ; the chlorine 
combines with the hydrogen to form hydrochloric acid, 
while the carbon is left free. 

The power to bleach is the most striking and useful prop- 
erty of chlorine. This property depends upon the fact 
that chlorine and water interact and ultimately liberate 
free oxygen ; the latter then decomposes the complex color- 
ing matter into colorless substances. If an envelope on 
which the postmark, or a pencil mark, is still visible is placed 
in moist chlorine, these marks will not be bleached because 
they are largely carbon; the writing ink will disappear. 
Litmus paper and many kinds of colored cloth are bleached 
by moist chlorine. The bleaching action of chlorine may be 
explained as follows : Chlorine and water form hypochlorous 
acid (HCIO), and then this very unstable acid decomposes 
according to the equation — 

HCIO = O + HC1 

Hypochlorous Oxygen Hydrochloric 
Acid Acid 



CHLORINE AND HYDROCHLORIC ACID 201 

Uses. — Chlorine gas is used extensively to manufacture 
bleaching powder. Liquid chlorine finds application in 
the manufacture of chlorides and in the extraction of gold 
from certain ores. Considerable chlorine gas is also used 
to bleach wood pulp; both gaseous and liquid chlorine are 
used to remove bromine from its compounds on a large 
scale. Bleaching powder is used as a germicide. 

Bleaching Powder is the main source of the chlorine used 
in the bleaching industries. It is sometimes called " bleach" 
or " chloride of lime." It is a yellowish white substance 
having a peculiar odor, which resembles that of chlorine. 
When dry, it is a powder, but on exposure to the air, it 
absorbs water and carbon dioxide, becomes lumpy and 
pasty, and loses some of its chlorine, owing to the formation 
and liberation of hypochlorous acid (HCIO). Acids like 
sulphuric and hydrochloric acid liberate from bleaching 
powder its " available chlorine," which varies from 30 to 
38 per cent in good qualities. The equations for the inter- 
action of acids and bleaching powder are usually written 
thus : — 

CaOCl 2 + H 2 S0 4 - 2 CI + CaS0 4 + H 2 

Bleaching Sulphuric Calcium 

Powder Acid Sulphate 

CaOCl 2 + 2HC1 = 2C1 + CaCl 2 + H 2 

Hydrochloric Calcium 

Acid Chloride 

The composition of bleaching powder has been much dis- 
cussed. The most reliable authority gives it the formula 
CaOCl 2 . When dissolved in water, bleaching powder forms 
calcium hypochlorite (Ca0 2 Cl 2 ) and calcium chloride (CaCl 2 ). 

Bleaching powder is manufactured by the interaction of chlorine 
gas and lime. Lime (calcium oxide, CaO) is carefully slaked with 
water to form calcium hydroxide (Ca(OH) 2 ). This powder is sifted 



202 INORGANIC CHEMISTRY 

into a large absorption chamber made of iron, lead, or tarred brick until 
the floor is covered with a layer three or four inches deep. The 
chlorine enters at the top and settles slowly to the floor. The simplest 
equation for the formation of bleaching powder may be written — 



Ca(OH) 2 


+ 2C1 = 


CaOCl 2 + 


H 2 


Calcium 


Chlorine 


Bleaching 


Water 


Hydroxide 




Powder 





Bleaching. — Immense quantities of bleaching powder are 
used to whiten cotton and linen goods and paper pulp. The 
pieces of cotton cloth as they come from the mill are 
sewed end to end in strips, which are stamped at the extreme 
ends with some indelible mark to distinguish each owner's 
cloth. These strips, which are often several miles long, 
are drawn by machinery into and out of numerous vats of 
liquors and water, between rollers, and through machines, 
until they are snow-white and ready to be finished (i.e. 
starched and ironed) or dyed. The whole operation re- 
quires three or four days. 

The preliminary treatment consists in singeing off the downy pile 
and loose threads by drawing the cloth over hot copper plates or 
through a series of gas flames. The object of the remaining opera- 
tions is threefold : (1) to wash out mechanical impurities, the fatty 
and resinous matter, and the excess of the different chemicals; (2) to 
remove matter insoluble in water; and (3) to oxidize the coloring 
matter by chlorine. The details of the process differ with the texture 
of the cloth and with its ultimate use. The threefold object above 
mentioned involves successively "liming," "souring/' "chemicking," 
and "souring," interspersed with frequent washing. The "liming" 
consists in boiling the cloth in a large kier, or vat, with lime, the 
"souring" in wetting it with dilute sulphuric or hydrochloric acid, 
and the "chemicking" in impregnating it with a weak solution of 
bleaching powder. Often the cloth is boiled at a certain stage with 
resin and sodium carbonate. The "liming" removes the resinous 
and the fatty matter, the first "souring" neutralizes traces of lime, 
and the second, which follows the "chemicking," liberates the chlorine 
in the fiber of the cloth. Frequent washing is absolutely necessary 
to remove the impure products of the chemical changes as well as 



CHLORINE AND HYDROCHLORIC ACID 



203 



the excess of lime and other alkali, acid, and chlorine. Should these 
be left, the cloth would be unevenly bleached, and its fiber would be 
weak. The cloth is finally treated with an antichlor, such as sodium 
hyposulphite, which removes the last traces of chlorine. 

Chlorides are formed when chlorine combines with other 
elements, and they are in general stable compounds. The 
simplest equations illustrating the combination of chlorine 
with certain elements are : — 



Na + 


Cl 


= NaCl 


Sodium 


Chlorine Sodium 
Chloride 


Sb + 


3C1 


= SbCl 8 


Antimony 




Antimony 
Trichloride 


Cu + 


2C1 


= CuClj 


Copper 




Copper 
Chloride 


p + 


3C1 


= PC1 8 


Phosphorus 




Phosphorus 
Trichloride 


H + 


Cl 


= HCl 


Hydrogen 




Hydrochloric 
Acid 



Chlorides are an important class of compounds, and they will 
be considered under the elements with which chlorine com- 
bines. (See also Chlorides, below.) 

Hydrochloric Acid 

Hydrochloric acid is the common name of a water solution 
of a very important compound of hydrogen and chlorine, 
viz. hydrogen chloride, HCl. Hydrogen chloride is a gas, 
which is very soluble in water. This solution is known in 
commerce as muriatic acid (from the Latin word muria, 
meaning brine), but it is more properly called hydrochloric 
acid. Hydrogen chloride is often called hydrochloric acid 
gas. 



204 INORGANIC CHEMISTRY 

Occurrence. — Hydrogen chloride occurs free in volcanic 
gases. The solution is one constituent of the gastric juice. 

Preparation. — The gas is prepared by heating concen- 
trated sulphuric acid and sodium chloride in an apparatus 
like that used for chlorine (Fig. 23). If the mixture is gently 
heated, the chemical change is represented thus : — 

NaCl + H2SO4 = HC1 + HNaS0 4 

Sodium Sulphuric Hydrochloric Acid Sodium 

Chloride Acid . Acid Sulphate 

But at a high temperature the equation for the reaction is — 
2 NaCl + H 2 S0 4 = 2 HC1 + Na 2 S0 4 

Sodium 
Sulphate 

The solution is prepared by passing the gas into water. 

If sulphuric acid is added to a solution of sodium chloride 
instead of the solid, little or no hydrogen chloride is liberated 
(unless heat is applied). This is due to the fact that the 
hydrogen chloride being very soluble in water remains in the 
sphere of action and tends to cause the equilibrium — 

NaCl + H2SO4 ^ HC1 + HNaS0 4 

It has already been pointed out that one way to displace equi- 
librium is to remove a product of the reaction. Thus, in this 
case the volatile hydrogen chloride escapes from the mixture 
of sulphuric acid and (solid) sodium chloride and allows the 
forward reaction to proceed. 

Commercial Hydrochloric Acid is manufactured by heat- 
ing a mixture of salt and sulphuric acid to a moderate 
temperature in a hemispherical cast iron retort, and conduct- 
ing the gas through an earthenware pipe into an absorbing 
tower ; the fused mass of acid sodium sulphate and salt is 
then subjected to a higher temperature, and the remainder 
of the gas is led into the absorbing tower. These towers are 



CHLORINE AND HYDROCHLORIC ACID 205 

high and filled with coke or pieces of brick over which water 
trickles ; as the hydrochloric acid gas passes up the tow T er, 
it is absorbed by the descending water, and concentrated 
acid flows from the bottom of the tower. The gas is usually 
cooled before it enters the towers. Sometimes the gas is 
conducted through huge earthenware jars before entering 
the towers. In these jars the gas and water are caused to 
flow constantly in opposite directions, thus insuring com- 
plete absorption. Hydrochloric acid can be manufactured 
synthetically, i.e. by burning hydrogen gas in chlorine 
gas. The equation for this reaction is — 
H + CI = HC1 

Hydrogen Chlorine Hydrochloric 
Acid 

Properties. — Hydrochloric acid gas is colorless. When 
it escapes into moist air, it forms fumes which are really 
minute drops of a solution of the gas in the moisture of the 
air. It has a choking, sharp taste, and irritates the lining 
of the nose and throat. The gas does not burn nor support 
combustion. It is about 1.25 times heavier than air, and 
may therefore be collected by displacement of air, like 
chlorine. One liter at 0° C. and 760 mm. weighs 1.641 gm. 
The critical temperature is +52° C, and the gas becomes 
a colorless liquid when subjected to a high pressure and 
moderately low r temperature. The extreme solubility of 
hydrochloric acid gas in water is one of its most striking 
properties. One liter of water will dissolve about 550 1. 
of gas, if both are at 0° C. and 760 mm. At the ordinary 
temperature about 500 1. of gas dissolve in 1 1. of water, 
and as the temperature rises, the solubility decreases. The 
solution is the familiar hydrochloric acid. The gas readily 
escapes, hence the acid forms fumes w T hen exposed to the 
air. Pure hydrochloric acid is a colorless liquid, but the 
commercial acid has a yellow color, usually due to iron 



206 INORGANIC CHEMISTRY 

compounds or to dissolved chlorine. Like the common 
acids, its solution reddens blue litmus and yields hydrogen 
by interaction with certain metals. In terms of the theory 
of electrolytic dissociation, hydrochloric acid is a strong 
acid; i.e. it dissociates to a considerable extent into ions, 
one kind being H + , the other being CI". 

The most concentrated acid contains about 40 per cent 
(by weight) of the compound (HC1), and its specific gravity 
is 1.2. When the concentrated acid is heated, the gas is 
evolved until the solution contains about 20 per cent of the 
acid, and then the liquid boils at 110° C. without further 
change in concentration. The dilute acid loses water until 
the same conditions prevail. 

Uses of Hydrochloric Acid. — Vast quantities are used 
to prepare the chlorine consumed in the manufacture of 
bleaching powder. Various chlorides are prepared from 
it, and it is one of the common acids used in the chemical 
laboratory and in many industries. 

Chlorides are formed by direct addition of chlorine to 
metals, as we have seen. They are also formed by the 
substitution of a metal for the hydrogen in hydrochloric 
acid. Chlorides, therefore, are salts of hydrochloric acid. 
They can be prepared in several ways, e.g. (1) by the inter- 
action of hydrochloric acid and metals, metallic oxides, 
or hydroxides, and (2) by the interaction of certain salts 
and hydrochloric acid or chlorides. The following equa- 
tions illustrate several of these methods : — 

Zn + 2HC1 = ZnCl 2 + 2H 

Zinc Zinc 

Chloride 

ZnO + 2HC1 = ZnCl 2 + H 2 

Zinc Zinc 

Oxide Chloride 






CHLORINE AND HYDROCHLORIC ACID 207 
Zn(OH) 2 + 2HC1 = ZnCl 2 + 2H 2 

Zinc Zinc 

Hydroxide Chloride 

AgNOa + HC1 = AgCl + HN0 8 

Silver Silver Nitric 

Nitrate Chloride Acid 

One molecule of a chloride may contain several atoms of 
chlorine. Often the name of the compound indicates this 
fact; e.g. manganese dichloride (MnCl 2 ), antimony tri- 
chloride (SbCl 8 ), phosphorus trichloride and pentachloride 
(PC1 3 and PCI5). If a metal forms two chlorides, the two 
are distinguished by modifying the name of the metal; 
the one containing the smaller proportions of chlorine ends 
in -ous, the one containing the larger ends in -ic. Thus, 
mercurous chloride is HgCl, but HgCl 2 is mercuric chloride. 
Similarly, we have ferrous chloride, FeCl 2 , and ferric chloride, 
FeCl 3 . 

The Test for Hydrochloric Acid and Chlorides. — Most 
chlorides are soluble in water, those of lead, silver, and 
mercury (-ous) being the only conspicuous exceptions. 
If silver nitrate is added to hydrochloric acid or to the 
solution of a chloride, a white, curdy precipitate of silver 
chloride is formed, w T hich (a) is insoluble in nitric acid, 
(6) soluble in warm ammonium hydroxide, and (c) turns 
purple in the sunlight. The invariable formation of silver 
chloride is the test for hydrochloric acid and soluble chlorides. 
This chemical change is a typical illustration of double de- 
composition. The equation for the chemical change is — 

HC1 + AgN0 3 = AgCl + HNO3 

Hydrochloric Silver Silver Nitric 

Acid Nitrate Chloride Acid 

By an inspection of this equation we see that both the hydro- 
chloric acid and silver nitrate decompose and the chemical" 
fragments, so to speak, recombine in a different way to 



208 INORGANIC CHEMISTRY 

form the two products. The same test is applicable to 
hydrochloric acid and a chloride, because both yield chlorine 
ions (Cl~). The test is for ionic chlorine, and the general 
ionic equation may be written thus : — 

C1- + Ag + ->AgCl 

Miscellaneous. — Besides hydrochloric acid, there are 
four other acids of chlorine ; they contain oxygen as well 
as hydrogen. These acids are hypochlorous acid (HCIO), 
chlorous acid (HC10 2 ), chloric acid (HC10 3 ), and perchloric 
acid (HCIO4). They are prepared with difficulty, and de- 
compose readily. Thus, hypochlorous acid is prepared 
by the interaction of dilute hydrochloric acid and sodium 
hypochlorite (NaCIO), but the procedure must be careful, 
because the resulting hypochlorous acid itself is decom- 
posed by hydrochloric acid. Hypochlorous acid also de- 
composes when its solution is warmed or exposed to the 
sunlight, oxygen gas and hydrochloric acid being the prod- 
ucts. This acid (HCIO) is formed to a slight extent when 
chlorine dissolves in water, thus : — 

2 CI + H 2 = HCIO + HC1 

Chlorine Water Hypochlorous Hydrochloric 

Acid Acid 

But the hypochlorous acid decomposes readily and under 
many conditions, especially when exposed to the sunlight; 
the equation for the reaction in the last-named case may be 
written : — 

HCIO = O + HC1 

Hypochlorous Oxygen Hydrochloric 
Acid Acid 

It is this free oxygen which is the effective agent in bleach- 
ing, though the unstable hypochlorous acid is the way 
station, so to speak, between the chlorine and the oxygen. 
The corresponding salts of two of the oxy-chlorine acids 



CHLORINE AND HYDROCHLORIC ACID 209 

are important, viz. the hypochlorites and chlorates. Solu- 
tions of potassium hypochlorite (Javelle's water) and sodium 
hypochlorite (Labarraque's solution) find application in the 
removal of stains from cotton and linen goods; " bleach 
liquors " consisting largely of hypochlorites are use as 
bleaching agents in some industries. Chlorates of potassium 
and sodium are used as a source of oxygen. 

Chlorine forms two oxides — chlorine monoxide (C1 2 0) 
and chlorine dioxide (C10 2 ). Each is an unstable yellowish 
brown gas. 

Problems and Exercises 

1. Calculate the volume of hydrochloric acid solution (having 
a specific gravity of 1.10 and containing 20.12 per cent of HC1) 
required to make 500 gm. of barium chloride from barium carbonate. 

2. What weight of chlorine can be prepared from 78 gm. of rock 
salt, containing 99 per cent of NaCl ? How much manganese diox- 
ide and sulphuric acid would be required? What volume would 
the chlorine occupy under standard conditions ? 

3. (a) Magnesium chloride heated in steam forms hydrogen 
chloride and magnesium oxide. Write the equation, (b) What 
volume of hydrogen chloride (at 0° C. and 760 mm.) can be formed 
from 50 gm. of magnesium chloride? 

4. (a) Calcium oxide and hydrochloric acid form calcium chlo- 
ride and water. Write the equation, (b) What volume of hydro- 
gen chloride (at 0° C. and 760 mm.) is needed to produce 50 gm. 
of calcium chloride? 

5. Write equations for the following reactions by applying the 
method outlined in Chapter VII : (a) Phosphorus and chlorine form 
phosphorus trichloride, (b) Phosphorus trichloride and chlorine 
form phosphorus pentachloride. (c) Aluminium and hydrochloric 
acid form hydrogen and aluminium chloride. 

6. Write ionic equations for the interaction of (a) hydrochloric 
acid and silver sulphate, and (6) calcium chloride and sodium sul- 
phate. 

7. Interpret the changes in 6 from the standpoint of solubility 
product. 

8. Interpret the preparation of hydrogen chloride from the 
standpoint of displacement of equilibrium. 



CHAPTER XIII 
Compounds of Nitrogen — Gay-Lussac's Law 

The most important compounds of nitrogen are ammonia 
(NH 3 ), nitric acid (HN0 3 ), and compounds related to them. 
Many animal and vegetable substances essential to life are 
compounds of nitrogen. 

Ammonia 

The term ammonia includes both the gas and its solution 
in water, though the latter is more accurately called am- 
monium hydroxide. 

Formation of Ammonia. — When vegetable and animal 
matter containing nitrogen decomposes or decays, the 
nitrogen and hydrogen are liberated in combination as 
ammonia. The odor of ammonia is often noticed near 
stables. If animal matter containing nitrogen is heated, 
ammonia is given off. The old custom of preparing ammonia 
by heating horns and hoofs in a closed vessel, i.e. by dry 
distillation, gave rise to the term " spirits of hartshorn." 
Soft coal contains compounds of nitrogen and of hydrogen, 
and when the coal is heated to make illuminating gas, one 
of the products is ammonia. 

Preparation. — Ammonia gas is prepared in the laboratory 
by heating ammonium chloride with a base, the mild base 
calcium hydroxide being usually used. The equation for 
the reaction is — 

2NH 4 C1 + Ca(OH) 2 = 2NH 4 OH + CaCl 2 

Ammonium Calcium Ammonium Calcium 

Chloride Hydroxide Hydroxide Chloride 

210 



COMPOUNDS OF NITROGEN 



211 



But the ammonium hydroxide is unstable and decomposes 
into ammonia gas and water, thus : — 



NH,OH 


= NH 3 + 


H 2 


Ammonium 


Ammonia 


Water 


Hydroxide 


Gas 





KJ=t 



JC 



The ammonia gas is very volatile and is usually collected 

by allowing it to flow into an inverted bottle and displace 

the air (Fig. 24) ; the term upward 

displacement is sometimes applied to 

this method of collecting a gas. The 

solution is prepared by conducting 

the gas into water. 

The main source of the ammonia ox com- 
merce is the ammoniacal liquor or gas liquor 
of the illuminating gas works. The gases 
which come from the retorts in which the 
coal is heated are passed into water, which 
absorbs the ammonia and certain other 
gases. This impure gas liquor is treated 
with an alkali to liberate the ammonia, which 
is absorbed in tanks containing hydrochloric —zj= 

acid or sulphuric acid. This solution upon FlG * 24 - Apparatus for 
the addition of an alkali (such as calcium 
hydroxide) gives up its ammonia, which is 
dissolved in distilled water, forming thereby 
the ammonium hydroxide or aqua ammoniae 
of commerce. 



The gas from the gen- 
erator flows through the 
delivery tube B to the 
top of the bottle D and 
displaces the air, which 
escapes through the hole 
in the block C 



Properties of Ammonia. — Ammonia 
gas is colorless. It has an exceed- 
ingly pungent odor, and if inhaled suddenly or in large 
quantities it brings tears to the eyes and may cause suf- 
focation. It is a light, volatile gas, being only .59 times 
as heavy as air. A liter of the gas at 0° C. and 760 mm. 
weighs about .77 gm. It will not burn in the air under 
ordinary conditions, nor will it support the combustion 



212 INORGANIC CHEMISTRY 

of a blazing stick ; but if the air is heated, or if considerable 
oxygen is mixed with the air, then a jet of ammonia gas 
may be made to burn (in a suitable apparatus) with a pale 
yellowish flame. 

Ammonia gas is easily liquefied, since its critical tem- 
perature is about +130° C; at 0° C. a pressure of about 
4.2 atmospheres causes liquefaction. At about —76° C. 
a white solid is produced. Liquefied ammonia is often 
called anhydrous ammonia, because it contains no water. 
It boils at about — 33.5° C. Hence, if it is exposed to 
the air or warmed in any way, it changes back to a gas, 
and in doing so absorbs considerable heat. This fact has 
led to the extensive use of liquid ammonia in the manu- 
facture of ice. 

When electric sparks are passed through ammonia gas, 
it decomposes into nitrogen and hydrogen. But if a mixture 
of nitrogen and hydrogen is sparked, these gases combine 
to form ammonia gas. In neither case, however, is the re- 
action complete. The final mixture always contains about 
98 per cent of nitrogen and hydrogen, and 2 per cent of 
ammonia gas. That is, the reaction is reversible and pro- 
ceeds until equilibrium is reached. The simplest equation 
for this reversible reaction is — 

NH 3 ^tN + 3H 

If water or acid is added during the sparking of the mixture 
of nitrogen and hydrogen, the ammonia is dissolved. Its 
removal displaces the equilibrium, and the reaction pro- 
ceeds to completion; i.e. all the nitrogen and hydrogen 
combine. 

Another marked property of ammonia gas is its solubility 
in water. A liter of water at 0° C. dissolves 1148 1. of gas 
(measured at 0° C. and 760 mm.), and at the ordinary tem- 
perature 1 1. of water dissolves about 700 1. of gas. This 



COMPOUNDS OF NITROGEN 213 

solution of the gas. is usually called ammonia, though other 
names are often applied to it; i.e. ammonium hydroxide, 
aqua ammonia?, or ammonia water. It gives off the gas freely, 
especially when heated, as may easily be discovered by the 
odor or by the formation of the dense white fumes of solid 
ammonium chloride (NH 4 C1) when the solution is exposed 
to hydrochloric acid. The volatility of ammonia was early 
detected, and the name volatile alkali w r as applied to it; 
the discoverer, Priestley, called the gas alkaline air. The 
gas can be completely removed from solution by boiling. 
The commercial solution is lighter than water (its specific 
gravity being about .88) and contains about 35 per cent 
(by weight) of the gas. Ammonium hydroxide has an alka- 
line reaction, neutralizes acids and forms salts, and acts 
in many respects like sodium hydroxide. In terms of the 
theory of electrolytic dissociation, ammonium hydroxide 
is a weak base ; i.e. it dissociates only to a slight degree (1.4 
per cent in N/10 solution at 18° C.) into ions (NHJ and OH"). 

Ammonium Compounds. — When ammonia gas is passed 
into water, some of the gas combines with the water and 
forms a solution of an unstable compound called ammonium 
hydroxide. Its formation may be represented thus : — 

NH 3 + H 2 = NH 4 OH 

Ammonia Water Ammonium 

Hydroxide 

Ammonia (NH 3 ) also unites directly with acids, thereby 
forming salts, thus : — 



NH 3 + HC1 = 


NH 4 C1 


Ammonia Hydrochloric 


Ammonium 


Gas Acid 


Chloride 



Analogous salts are ammonium sulphate and ammonium 
nitrate. 



214 INORGANIC CHEMISTRY 

Ammonium hydroxide is a base. It neutralizes acids and 
forms salts, thus : — 

NH 4 OH + HC1 = NH 4 C1 + H 2 

Ammonium 
Chloride 

2NH 4 OH + H 2 S0 4 = (NH 4 ) 2 S0 4 + 2H s O 

Ammonium 
Sulphate 

These salts, ammonium chloride and ammonium sulphate, 
are strictly analogous to sodium salts. Thus, we have — 

Sodium Salts Ammonium Salts 

NaCl NH 4 C1 

NaN0 3 NH 4 N0 3 

Na 2 S0 4 (NH 4 ) 2 S0 4 
etc. etc. 

Hence, it is believed that ammonium compounds contain 
a group of atoms which acts chemically like an atom of a 
metal. This group of atoms is called ammonium, and its 
formula is NH 4 . Ammonium has never been separated 
from its compounds, or if it has, it is so unstable that it 
immediately decomposes into ammonia gas and hydrogen. 
So also ammonium hydroxide has never been obtained free, 
for it decomposes readily into ammonia gas and water. How- 
ever the properties of ammonium hydroxide leave no doubt 
but that it is a compound of hydroxyl and ammonium. 
Ammonium is called a radical, because it is the root of a 
series of compounds. 

Uses of Ammonium Hydroxide. — Ammonium hydroxide 
is widely used as a cleansing agent (especially for the removal 
of grease), as a restorative in case of fainting or of inhaling 
irritating gases, in dyeing, and calico printing, and in the 
manufacture of dyestuffs, sodium bicarbonate, and am- 
monium compounds. 



COMPOUNDS OF NITROGEN 



215 



The Use of Ammonia as a Refrigerant and in making Ice 
depends on the fact that liquefied ammonia (not ammonia 
water) changes rapidly into a gas and thereby absorbs heat. 
Hence, if liquefied ammonia is allowed to flow through a pipe 
immersed in a solution of sodium chloride or calcium chloride 
(technically called a brine), the ammonia evaporates in the 
pipe and cools the brine, which may be used directly as a re- 
frigerant or for making ice. In some cold storage plants, 
breweries, packing houses, and sugar refineries this cold brine 
is circulated through a system of pipes placed in the rooms 
where a low temperature is desired. 



Cold "Water trickling over' the 

ammonia pipes to condense 

the compressed gas 




^ 



Ammonia pump 
1 1 Low Pressure ^f High Pressure 



Expansion valve 
Brine pump 

Fig. 24 a. — Diagram of an ice-making plant. 



In the operation of an ice-making plant (Fig. 24 a) liquefied 
ammonia is forced from a tank into a series of pipes which are 
submerged in an immense vat filled with brine. Large galvanized 
iron cans containing pure water to be frozen are immersed in the 
brine, which is kept below the freezing point of water by rapid 
evaporation of the ammonia in the pipes. After several hours the 
water in the cans is frozen into ice. As fast as the ammonia gas 
forms in the pipes, it is removed by exhaust pumps into another tank, 
where it is recondensed to liquefied ammonia and conducted, as 
needed, into the first tank ready for renewed use. The ammonia is 
thus used over and over without appreciable loss. The pure water 
is often obtained from an artesian well. Most ocean steamers have 
an ice plant and cold storage room ; in cities in warm climates 
manufactured ice is a common commodity. 



216 



INORGANIC CHEMISTRY 



Composition of Ammonia Gas. — Many facts show that ammonia 

gas is composed of nitrogen and hydrogen in the ratio of 1 to 3 by 

volume and 14 to 3 by weight. The volumetric com- 

fl±l position is shown in two ways. The first depends upon 

the fact that ammonia and chlorine interact thus : — 

2NH 3 + 3C1 2 = N 2 + 6HC1 

Ammonia Chlorine Nitrogen Hydrochloric 

Gas Acid 

A tube (Fig. 25) filled with a known volume of 
chlorine is provided with a funnel through which con- 
centrated ammonium hydroxide is dropped into the 
chlorine, until the reaction ceases. After the excess of 
ammonia is neutralized with sulphuric acid, the volume 
of nitrogen left is found to be one third of the original 
volume of chlorine gas. Now hydrogen and chlorine 
combine in equal volumes; hence, the volume of hydro- 
gen withdrawn from the added ammonia must be 
Fig. 25. — Ap- equal to the original volume of chlorine. But this 
paratus for volume is three times the volume of nitrogen ; therefore 
determining there is three times as much hydrogen as nitrogen by 
. ? com P°- volume in ammonia gas. 

The gravimetric composition of ammonia gas is 
found by oxidizing it, and weighing the water and 
nitrogen, which are the only products. The results show that four- 
teen parts of nitrogen combine with three parts of hydrogen. 

Nitric Acid 

Nitric Acid is one of the most useful compounds of nitro- 
gen. It was known to the alchemists, who used it to pre- 
pare a mixture which dissolves gold. (See Aqua Regia, below.) 
Nitric acid is used in the preparation of many nitrogen com- 
pounds. 

Formation of Nitric Acid. — When moist animal or vege- 
table matter containing nitrogen decays in the presence 
of an alkali, nitric acid is formed ; but it is neutralized at 
once by the alkali, so nitrates — salts of nitric acid — are the 
final products. This chemical change is known as nitri- 



sition of am- 
monia gas. 



COMPOUNDS OF NITROGEN 217 

fication, and it is caused, or largely influenced, by minute 
living organisms called bacteria. 

Nitric acid is formed when electric sparks are passed 
through moist air. The nitrogen and oxygen combine to 
some extent and form nitric oxide (NO), which unites with 
oxygen and forms nitrogen dioxide (N0 2 ). This gas and 
water form nitric acid thus : J — 

3 N0 2 + H 2 = 2 HN0 3 + NO 

Nitrogen Water Nitric Nitric 

Dioxide Acid Oxide 

This chemical change is now being applied on a commercial 
scale. Air is forced through a tube in which a powerful 
electric arc is spread out into a disk by an electromagnet. 
The nitrogen dioxide is absorbed in water or in a mixture of 
water and lime, thereby forming nitric acid or calcium nitrate. 

Preparation. — Nitric acid is prepared in the laboratory 
by heating concentrated sulphuric acid with a nitrate 
in a glass retort ; the nitric acid distills into a receiver, which 
is kept cool by water. The reaction at ordinary temperatures 
is represented by the equation — 

NaN0 3 + H2SO4 ^t HNO3 + HNaS0 4 

Sodium Sulphuric Nitric Acid Sodium 

Nitrate Acid Acid Sulphate 

The reaction is reversible and equilibrium is soon established. 
But since nitric acid boils at 86° C, it is removed by gentle 
heat and recovered by condensation. Its removal displaces 
the equilibrium and allows the forward reaction to proceed 
to completion. At a high temperature sodium sulphate 
(Na 2 S0 4 ) is formed, but since part of the nitric acid is de- 
composed, excessive heat is avoided. 

Nitric acid is manufactured by heating sodium nitrate 
and sulphuric acid in a large cast-iron retort and condensing 



218 INORGANIC CHEMISTRY 

the vapors in huge glass or earthenware bottles; the last 
bottle is connected with a tower filled with coke over which 
water trickles to absorb the vapors which escape from the 
bottles. The acid vapors are also often condensed in earthen- 
ware or glass tubes. 

Properties. — Pure nitric acid is a colorless liquid, but the 
commercial acid is yellow or reddish, owing to absorbed 
nitrogen compounds, chlorine, or iron compounds. It 
decomposes slowly in the sunlight or when heated, and 
a brownish gas may often be seen in bottles of nitric acid. 
It absorbs water, and forms irritating fumes when exposed 
to the air. The specific gravity of the commercial acid is 
about 1.42, and it contains approximately 70 per cent of the 
real acid (HN0 3 ), the rest being water. If the water is 
removed by slowly distilling the commercial acid with con- 
centrated sulphuric acid, the product contains from 94 to 
99 per cent of the real acid, and its specific gravity is about 
1.51. When nitric acid is boiled, it loses either acid or water 
until the liquid contains approximately 68 per cent of nitric 
acid, and then it continues to boil unchanged in concentration 
at 120° C. (See Hydrochloric Acid.) 

A solution of nitric acid has the properties of acids to 
a marked degree. In terms of the theory of electrolytic 
dissociation it is one of the strongest acids ; i.e. it dissociates 
largely into ions, one kind being H + , the other being N0 3 ~. 

Nitric acid is very corrosive. It turns protein, e.g. the skin, 
a yellow color owing to the formation of xanthoprotein ; the 
concentrated acid causes serious burns and should be used 
with extreme care. Nitric acid decomposes readily, espe- 
cially when hot, and is therefore an energetic oxidizing agent. 
Glowing charcoal continues to burn in the acid, while straw, 
sawdust, hair, and similar substances are charred and even 
inflamed by it. Iron sulphide (FeS) heated with nitric acid 



COMPOUNDS OF NITROGEN 219 

becomes oxidized to iron sulphate (FeS0 4 ). It interacts 
readily and often violently with metals, metallic oxides, 
and hydroxides, forming a variety of products, especially 
nitrates. 

Uses of Nitric Acid. — Nitric acid is one of the common 
laboratory acids. Large quantities are used in the manu- 
facture of nitrates, dyestuffs, sulphuric acid, nitroglycerin, 
and guncotton, and in the refining of gold and silver. 

Nitrates. — Nitric acid is monobasic and forms a series of 
well-defined salts called nitrates. The interaction of nitric 
acid and most metals is exceedingly vigorous, and for this 
reason, probably, the alchemists called the acid aqua fortis 
— strong water. The reaction varies with the metal, con- 
centration of the acid, temperature, and the presence of 
resulting compounds ; one product is usually a soluble nitrate, 
though some metals, such as tin and antimony, form in- 
soluble oxides. The gaseous products are usually oxides 
of nitrogen, especially colorless nitric oxide (NO), which 
quickly forms brown nitrogen dioxide (N0 2 ) in the air. Hy- 
drogen is seldom liberated so that it can be collected; it 
generally reduces the nitric acid to nitric oxide (NO) and 
water. Ammonia gas (NH 3 ) and even nitrogen are some- 
times formed. The reaction between moderately dilute 
nitric acid and copper is typical of some metals and may be 
written thus : — 



3 Cu + 8 HN0 3 = 


= 3 Cu(N0 3 ) 2 + 2 NO + 4 H 2 


Copper Nitric 


Copper Nitric Water 


Acid 


Nitrate Oxide 



This equation is really made up of three equations : — 

(1) 2 HN0 3 = 3 O + 2 NO + H 2 

(2) 3 Cu + 3 = 3 CuO 



Copper 
Oxide 



(3) 3 CuO + 6 HNO3 = 3 Cu(N0 3 ) 2 + 3 H 2 



220 INORGANIC CHEMISTRY 

Eliminating 3 O from (1) and (2) and 3 CuO from (2) and 
(3), the remaining terms make up the complete equation. 
In the case of zinc, which is typical of other metals, the 
equations are : — 

(1) 3 Zn + 6 HN0 3 = 3 Zn(N0 3 ) 2 + 6 H 

(2) 6 H + 2 HNO3 = 2 NO +4 H 2 

Eliminating the common factor (6 H), the complete equation 
is : — 

3 Zn + 8 HNO3 = 3 Zn(N0 3 ) 2 + 2 NO + 4 H 2 

Nitrates, as a rule, are very soluble in water. They behave 
in various ways when heated. Some, like sodium and potas- 
sium nitrates, lose oxygen and pass into nitrites ; others, like 
copper nitrate, form an oxide of the metal, an oxide of nitro- 
gen, and oxygen ; and one, ammonium nitrate, decomposes 
into water and nitrous oxide (N 2 0). Since many nitrates, 
when heated, give up oxygen, they are powerful oxidizing 
agents. Potassium nitrate dropped on hot charcoal burns 
the charcoal vigorously and rapidly. This kind of chemical 
action is called deflagration. Nitrates have numerous uses, 
and these, as well as their special properties, will be treated 
under their respective metals. 

The Test for Nitrates (and of course for nitric acid) is as 
follows : Add to the nitric acid or the solution of the nitrate 
an equal volume of concentrated sulphuric acid, and cool the 
mixture. Upon the cool mixture pour carefully a cold, di- 
lute, freshly prepared solution of ferrous sulphate. A dark 
brown layer appears where the two liquids meet, owing 
to the formation of an unstable compound which has ap- 
proximately the composition represented by 3 FeS0 4 . 2 NO. 
Owing to the solubility of nitrates, the nitrate ion (N0 3 ~) 
cannot be precipitated. 



COMPOUNDS OF NITROGEN 221 

Nitrous Acid, HN0 2 , is not easily obtained in the free state 
owing to its instability, but the nitrites are stable compounds. 
Potassium nitrite (KN0 2 ) and sodium nitrite (NaN0 2 ) are 
formed by removing the oxygen from the corresponding 
nitrate by heating with lead. Nitrites give a brown mix- 
ture of nitric oxide and nitrogen dioxide when treated with 
sulphuric acid, and are thus readily distinguished from 
nitrates. 

Oxides of Nitrogen 

There are five oxides of nitrogen : — 



Name 


Formula 


Nitrous Oxide 


N 2 


Nitric Oxide 


NO 


Nitrogen Trioxide 


N 2 o 3 

N0 2 


Nitrogen Dioxide 


Nitrogen Pentoxide 


N 2 5 







Only three of these are important ; viz. nitrous oxide, nitric 
oxide, and nitrogen dioxide. 

Nitrous Oxide, N 2 0, is one of the numerous decomposition 
products of nitric acid, but it is usually prepared by decom- 
posing ammonium nitrate. This salt, if gently heated, first 
melts and then decomposes into nitrous oxide and water; 
the gas can be collected over water, preferably warm water. 
The equation for the chemical change is — 



NH4N03 


N 2 


+ 


2H 2 


Ammonium 


Nitrous 




Water 


Nitrate 


Oxide 







This colorless gas has a sweet taste and a faint but pleasant 
odor. It is less soluble in hot than in cold water. The 



222 INORGANIC CHEMISTRY 

gas does not burn, but it supports the combustion of many 
burning substances, though not so vigorously as oxygen does. 
Sulphur, for example, will not burn in nitrous oxide, unless 
the sulphur is hot and well ignited at first ; very fine iron 
wire, if well ignited, burns in the gas, but the combustion is 
not so conspicuous as in oxygen. The products of the chem- 
ical change are oxides and nitrogen. The most striking 
property of nitrous oxide is its effect on the human system. 
If breathed for a short time, it causes more or less nervous 
excitement, sometimes manifested by laughter. The gas 
was called " laughing gas" by Davy. If breathed in large 
quantities, it slowly produces unconsciousness. The gas is 
often administered when unconsciousness is desired for a 
short time, as in dentistry. It is easily liquefied by cold, by 
pressure, or by both together, since its critical temperature 
is about +37° C. It is often used in liquid form to furnish 
the gas itself and to produce low temperatures. 

Nitrous oxide was discovered by Priestley in 1776; but its com- 
position was not explained until 1799, when Davy, by an extensive 
study of its properties, proved it to be an oxide of nitrogen. In his 
enthusiasm Davy wrote a friend: " This gas raised my pulse upward of 
twenty strokes, made me dance about the laboratory as a madman, 
and has kept my spirits in a glow ever since." It is needless to say 
that the usual results are more quieting. 

Nitric Oxide, NO, has long been known, since it is the 
usual gaseous product formed by the interaction of nitric 
acid and metals. It is conveniently prepared by the inter- 
action of copper and dilute nitric acid (sp. gr. 1.2). The 
equation for the complex chemical change is written thus : — 

3Cu + 8HN0 3 = 2 NO + 3Cu(N0 3 ) 2 + 4H 2 

Copper Nitric Nitric Copper 

Acid Oxide Nitrate 

The gas thus prepared is impure, and it is customary to use 
ferrous sulphate and nitric acid as a source of the pure gas. 



COMPOUNDS OF NITROGEN 223 

Nitric oxide is a colorless gas, but upon exposure to the air 
it combines at once with oxygen, forming dense reddish 
brown fumes of nitrogen dioxide. The equation for this 
change is — 

2 NO + 2 = 2 N0 2 

Nitric Nitrogen 

Oxide Dioxide 

This property distinguishes nitric oxide from all other gases. 
It does not burn, nor does it support combustion unless the 
burning substance (e.g. phosphorus or sodium) introduced 
is hot enough to decompose the gas into nitrogen and oxy- 
gen, and then, of course, the liberated oxygen supports the 
combustion. It is only slightly soluble in water. 

Nitrogen Dioxide (or Peroxide), N0 2 , is the reddish gas 
formed by the union of nitric oxide and oxygen. Thus : — 

2 NO + 2 = 2N0 2 

Nitric Nitrogen 

Oxide Dioxide 

It is also produced by heating certain nitrates. Thus : — 
Pb(N0 3 ) 2 = 2N0 2 + PbO + O 

Lead Nitrogen Lead Oxygen 

Nitrate Dioxide Oxide 

The fumes of nitrogen dioxide usually appear when nitric 
acid and metals interact, but, as already stated, the fumes 
are not produced at first, being the result of a second chemical 
change when the nitric oxide combines with oxygen of the 
air. 

Nitrogen dioxide has a disagreeable odor, and if breathed 
in moderately large quantities, it is poisonous. It interacts 
with water and under ordinary conditions jdelds nitric acid 
and nitric oxide. The gas also dissolves in concentrated 
nitric acid, forming fuming nitric acid, which is a powerful 
oxidizing agent. 



224 INORGANIC CHEMISTRY 

When the gas called nitrogen dioxide is sealed in a tube and 
cooled, the color changes from red-brown to pale yellow ; a yellow 
liquid forms at about 26° C. and a nearly colorless solid at about 
— 12° C. The yellow substance is nitrogen tetroxide (N 2 4 ). 
Upon heating the tube the red-brown gas reappears ; at about 
150° C. it is dark red-brown and is nitrogen dioxide (N0 2 ). These 
changes show that at ordinary temperatures the tube contains a 
mixture of the two gases which are in equilibrium, thus : — 



N 2 4 


<— 


2N0 2 


Nitrogen 




Nitrogen 


Tetroxide 




Dioxide 



Although the red-brown gas as ordinarily seen is a mixture, it is 
called nitrogen dioxide. The formation of nitrogen dioxide in- 
volves absorption of heat. Hence increased temperature favors 
the forward reaction (see Le Chatelier's law). 

Nitrogen Trioxide, N2O3, and Nitrogen Pentoxide, N 2 5 , are unstable 
compounds and have no practical importance. They are the anhy- 
drides of nitrous and nitric acids. 

Aqua Regia is an old term used by the alchemists and still 
applied to a mixture of concentrated nitric and hydrochloric 
acids (1 vol. to 3 vol.). The expression means " royal water/' 
and indicates that the mixture dissolves gold (a " noble " 
metal), which is insoluble in either acid alone. Its solvent 
power depends mainly upon the free chlorine which is pro- 
duced in the mixture by the oxidizing action of the nitric 
acid, thus : — 

HNO3 + 3 HC1 = 2 CI + NOC1 + 2 H 2 

Aqua regia reacts energetically with metals, and the product 
of the reaction is always a chloride of the metal. 

Gay-Lussac's Law of Gas Volumes 

Several gaseous compounds of nitrogen illustrate Gay- 
Lussac's law of gas volumes. Experiment shows the follow- 
ing facts about these gases and certain others previously 
studied : — 



COMPOUNDS 
Combination of 



OF NITROGEN 
Gases by Volume 



225 



Volumes of Combining Gases 


Volumes of Gaseous Peoducts 


2 vol. hydrogen 


2 vol. water vapor 


1 vol. oxygen 




1 vol. hydrogen 


2 vol. hydrochloric acid gas 


1 vol. chlorine 




3 vol. hydrogen 


2 vol. ammonia gas 


1 vol. nitrogen 




2 vol. nitrogen 


2 vol. nitrous oxide gas 


1 vol. oxygen 




1 vol. nitrogen 


2 vol. nitric oxide gas 


1 vol. oxygen 




1 vol. nitrogen 


2 vol. nitrogen dioxide gas 


2 vol. oxygen 




2 vol. nitrogen 


2 vol. nitrogen trioxide gas 


3 vol. oxygen 





It is clear from the above table that small whole numbers 
express the relation existing between the volumes of the com- 
bining gases and the volume of the gaseous product. This 
simple relation is general and was summarized in 1808 by 
the French chemist Gay-Lussac in the form of a law r , thus : — 

Gases combine in volumes which bear a simple numerical 
ratio to each other and to the volume of their gaseous product. 

Additional illustrations of this fundamental law will be 
given in subsequent chapters. (See especially Chapters XV 
and XVI (oxides of carbon and hydrocarbons.) See also 
gas equations, Chapter XIV.) 

Problems 

1. How many grams of ammonia gas can be obtained from 2140 
gm. of ammonium chloride by heating with lime ? 

2. Calculate the percentage composition of (a) ammonium chlo- 



226 INORGANIC CHEMISTRY 

ride, (6) ammonium hydroxide, (c) ammonium sulphate, (d) am- 
monium nitrate. 

3. What weight of pure sodium nitrate is needed to produce a 
metric ton of pure nitric acid ? 

4. How many grams of each element in 27 gm. of pure nitric 
acid? 

5. What weight of pure nitric acid can be obtained from a metric 
ton of sodium nitrate (95 per cent pure)? 

6. Will sodium nitrate or potassium nitrate yield the greater 
weight of nitric acid? 

7. What volume of sulphuric acid solution having a specific 
gravity of 1.8354 and containing 93.19 per cent of H 2 S0 4 is needed 
to convert 10 kg. of pure sodium nitrate into pure nitric acid? 

8. What volume of nitric acid solution having a specific gravity 
of 1.4 and containing 65.3 per cent of HN0 3 can be obtained by the 
interaction of sulphuric acid having the concentration given in 
Problem 7 and 25 metric tons of pure sodium nitrate? 

9. What volume of oxygen is needed to combine with the 
hydrogen obtained by passing electric sparks through 150 cc. of 
ammonia gas until equilibrium is reached? 

10. What weight and what volume of nitrous oxide can be 
prepared from 72 gm. of ammonium nitrate ? (Standard conditions. ) 

11. What weight of nitric oxide is formed by the interaction of 
nitric acid and 45 gm. of copper? What weight of nitrogen dioxide 
will the nitric oxide form? 

12. What weight and what volume of oxygen will be needed to 
convert (a) 70 gm. and (b) 70 1. of nitric oxide into nitrogen dioxide? 
(Standard conditions.) 

13. (a) What volume of oxygen is needed to convert 10 1. of 
NO into N0 2 ? (b) What volume of nitrogen, to convert 12 1. of 
hydrogen into NH 3 ? 

14. Calculate the volume of gas formed in each of the following 
reactions : (a) 150 cc. of hydrogen and sufficient nitrogen ; (b) hy- 
drogen and 150 cc. of nitrogen; (c) 100 cc. of oxygen and sufficient 
nitric oxide. 

15. Calculate the formulas corresponding to (a) N = 46.666, 
0-53.333; (6)N = 40, = 45.71, H = 14.28; (c)N = 35, = 60, 
H = 5. What is the name of each compound ? 



CHAPTER XIV 

Atomic and Molecular Weights — Valence 

Extended application of atomic and molecular weights has- 
been made in the foregoing pages. In the present chapter 
we shall consider the methods by which both atomic and 
molecular weights are determined; several cognate prin- 
ciples of fundamental importance will also be discussed. 

Retrospect. — Before beginning the development of atomic 
and molecular weights it will be helpful to review certain facts 
and assumptions which bear directly upon this subject. It 
will be recalled that chemical compounds have a definite com- 
position by weight, and furthermore that if the proportions 
of the elements in a series of compounds containing the same 
elements are expressed in a special way, the composition is 
revealed as a simple multiple relation. Again, it will be re- 
membered that gases exhibit a striking similarity of behav- 
ior not only when they are subjected to heat and pressure 
but also when they undergo chemical transformations. 
These phenomena are summarized in the laws of Boyle, 
Charles, and Gay-Lussac, which have already been discussed. 
It has also been seen that in chemical changes neither energy 
nor matter is lost; transformations occur, but both energy 
and matter are conserved. Finally we should not forget that 
the atomic theory offers an acceptable explanation of the com- 
position of matter and of certain aspects of chemical change 
by assuming that atoms are the gravimetric units of chemical 
change and by their combinations form molecules which in 
turn decompose wholly or in part, thereby producing the 
varied and complicated phenomena succinctly called chem- 
ical changes. These atoms have a relative and unvarying 

227 



228 INORGANIC CHEMISTRY 

weight called the atomic weight — alike for each atom of the 
same element, unlike for each atom of different elements. So 
also each molecule of a compound has a relative and unvary- 
ing weight called the molecular weight, which is the sum of 
the weights of the atoms in one molecule of the compound. 

Our present task is to discuss the methods of determining 
these relative weights of atoms and molecules, and to sup- 
plement the data collected in this brief retrospect by certain 
laws, principles, and theories. 

Determination of Atomic Weights. — The atomic weight 
of an element, as already stated, is a relative weight. It is 
a number expressing the relation of the weight of the atom 
of a given element to the weight- of the atom of some element 
chosen as a standard. Thus, if we say the atomic weight of 
nitrogen is 14, we mean that the ratio of the weight of the 
nitrogen atom and the weight of the hydrogen atom is 14 to 
1, provided we adopt 1 as the weight of the hydrogen atom; 
or we mean that the ratio of the weight of the nitrogen atom 
and the weight of the oxygen atom is 14 to 16, provided we 
adopt 16 as the weight of the oxygen atom. Hydrogen was 
the standard for many years. However, since oxygen com- 
bines readily with a large number of elements and forms 
compounds which are quite suitable for experimental work, 
the oxygen atom has been adopted as the international 
standard atom and the weight 16 has been given to it. 

When compounds are analyzed, the results show the pro- 
portions by weight in which the constituent elements are 
combined. If one molecule of a compound contained only 
one atom each of the united elements, the relative weights 
of the atoms could easily be determined. Thus, approxi- 
mately 8 parts of oxygen combine with 1 part by weight of 
hydrogen to form 9 parts of water; assuming that a mole- 
cule of water contains only 1 atom of each element, it is 



ATOMIC AND MOLECULAR WEIGHTS — VALENCE 229 

evident that the atomic weight of hydrogen would be 2 if 
oxygen is 16. But before accepting this conclusion we 
should not overlook the existence of another compound of 
hydrogen and oxygen (called hydrogen peroxide) in which 
1 part by weight of hydrogen combines with 16 parts of 
oxygen. Assuming as before that a molecule of hydrogen 
peroxide contains 1 atom of each element, the conclusion is 
forced upon us that the atomic weight of hydrogen would be 
1 if oxygen is 16. But the atomic weight of hydrogen can- 
not be both 1 and 2 ! If we only knew which of the com- 
pounds contained a single atom of each element to the 
molecule, — granting for the time being that such is the 
case, — the problem would be simple. Analysis does not 
reveal the number of atoms in a molecule. Obviously addi- 
tional data are indispensable. 

The determination and final selection of the atomic weight 
of an element is based upon (1) the equivalent weight of the 
element, (2) the molecular weights of several compounds of 
the element, and (3) accurate chemical analysis of selected 
compounds of the element. 

Equivalent Weights. — The chemical analysis of a com- 
pound, as already stated, gives the proportion by weight 
in which the constituents are combined. If these propor- 
tions are stated in a special way instead of the customary 
form of percentage, certain important relations are revealed. 
Thus, analysis of ferrous oxide yields approximately 77.77 per 
cent of iron and 22.22 per cent of oxygen; but if we substitute 
8 for 22.22, the weight of iron becomes 28, because 22.22 and 
77.77 are in the same ratio as 8 and 28. When a similar 
process of modification is applied to the percentage of the 
elements in different compounds, a series of numbers is ob- 
tained known as the equivalent weights of the elements. 
By definition, the equivalent weight of an element is the 



230 



INORGANIC CHEMISTRY 



number of grams which combines with or displaces 8 gm. of 
oxygen. A partial summary of numerous experiments gives 
the following : — 

Table op Equivalent Weights 



Element 


Equivalent Weight 


Element 


Equivalent Weight 


Oxygen .... 


8 


Iron (-ous) . 


27.92 


Aluminium . . 


9.03 


Iron (-ic) . . 


18.62 


Bromine .... 


79.92 


Magnesium . 


12.16 


Calcium .... 


20.03 


Mercury (-ic) 


100.3 


Carbon .... 


3 


Potassium . 


39.10 


Chlorine .... 


35.46 


Silver .... 


107.88 


Copper (-ous) . 


63.57 


Sodium . . . 


23 


Copper (-ic) . . 


31.79 


Sulphur . . . 


16.03 


Hydrogen . . . 


1.008 


Zinc .... 


32,68 



This list might be extended to include all the elements 
which form compounds. These numbers are sometimes 
called combining numbers, combining weights, or simply 
equivalents. The term equivalent weights is preferable, 
because they actually are the weights chemically equivalent 
to each other. Thus, if we start with hydrogen chloride 
(HC1), 1 gm. of hydrogen — to take a convenient denomi- 
nation — is combined with 35.46 gm. of chlorine, and this 
gram of hydrogen can be replaced chemically by 32.68 gm. 
of zinc, 12.16 gm. of magnesium, 39.10 gm. of potassium, 
23 gm. of sodium, and so on. These elements are chem- 
ically equivalent in the ratio of these weights. 

Equivalent weights are readily found by experiment in most cases. 
Various methods are used, and the equivalent weight is not always 
found directly in terms of oxygen. The equivalent of hydrogen is 
found by passing hydrogen over hot copper oxide — as in the deter- 
mination of the gravimetric composition of water. The equivalent 
of magnesium is found by filling a graduated tube with dilute hydro- 
chloric acid, inserting the magnesium, inverting the tube in a dish, 
and measuring the volume of the liberated hydrogen; knowing the 
weight of a liter of hydrogen, the weight of the liberated hydrogen 



ATOMIC AND MOLECULAR WEIGHTS — VALENCE 231 

can be calculated, and from its weight the weight of magnesium 
equivalent to 1 gm. of hydrogen can be found by proportion. The 
equivalent weight of magnesium can also be found by heating a known 
weight of magnesium in the air — taking care, of course, not to lose 
the product. Equivalent weights of copper and other metals as well 
as of sulphur and of carbon can be found by passing oxygen over 
these substances in a tube which contains or is attached to a suitable 
apparatus for retaining the product. The interaction of metals and 
acids provides a simple method of finding the equivalent of zinc, 
aluminium, and iron; the interaction of water with sodium and 
with calcium permits the determination of the equivalent weights of 
these metals; and the displacement of metals from solutions of their 
compounds by such metals as zinc and magnesium provides another 
general method. 

Equivalent weights have been very carefully determined 
by experiment. Comparison of the equivalent weight with 
the atomic weight of the same element reveals an important 
relation, as may be seen by the following : — 



Table of Equivalent Weights and Atomic Weights 



Element 



Oxygen . . 
Aluminium . 
Bromine 
Calcium . . 
Carbon . . 
Chlorine . . 
Copper (-ous) 
Copper (-ic) 
Hydrogen . 
Iron (-ous) . 
Iron (-ic) . 
Magnesium 
Mercury (-ic) 
Potassium . 
Silver . . 
Sodium . . 
Sulphur . . 
Zinc . . . 



Equivalent 
Weight 


Atomic Weight 


Multiple 


8 


16 


2 


9.03 


27.1 


3 


79.92 


79.92 


1 


20.03 


40.06 


2 


3 


12 


4 


35.46 


35.46 


1 


63.57 


63.57 


1 


31.79 


63.57 


2 


1.008 


1.008 


1 


27.92 


55.85 


2 


18.62 


55.85 


3 


12.16 


24.32 


2 


100.3 


200.6 


2 


39.10 


39.10 


1 


107.88 


107.88 


1 


23 


23 


1 


16.03 


32.07 


2 


32.68 


65.37 


2 



232 INORGANIC CHEMISTRY 

An examination of this comparative table shows an integral 
relation between equivalent weights and atomic weights. 
In other words, the atomic weight of an element is identical 
with its equivalent weight or is a simple integral multiple 
of it. The significance of this relation will be discussed after 
the subject of molecular weights has been considered. 

Determination of Molecular Weights. — Before describing 
the actual methods . employed in determining molecular 
weights, it will be necessary to discuss the kinetic theory of 
gases and Avogadro's hypothesis. These theoretical prin- 
ciples underlie the interpretation of the results obtained by 
experiment and assist in the correlation of the properties of 
gases summarized by the laws of Boyle, Charles, and Gay- 
Lussac. 

Kinetic Theory of Gases and Avogadro's Hypothesis. — 

Extensive study of gases shows that in general they conform 
to fundamental laws. These laws indicate the uniform and 
simple structure of all gases. The theory proposed to ex- 
plain the uniform behavior of gases* is called the kinetic 
theory. According to this theory gases are conceived to 
consist of molecules, moving constantly and rapidly in all 
directions; these particles are also conceived to have perfect 
elasticity, and in their movements in the space which is large 
compared with their own bulk they collide with each other 
or with the walls of the containing vessel, rebound, and con- 
tinue to move without loss of energy; furthermore, the mole- 
cules of a gas move in straight lines and have little or no 
tendency to repel or adhere to each other; i.e. they are inde- 
pendent particles separated by an average distance much 
greater than their own diameters. Recasting this theory 
into a concrete form, a vessel of oxygen gas contains a vast 
number of molecules, flying about rapidly in the space, 
striking each other and the walls of the vessel, rebounding 



ATOMIC AND MOLECULAR WEIGHTS — VALENCE 233 

after each collision, some, eventually, all, flying out and min- 
gling with the air, but not combining with each other or com- 
pressing each other, or even adhering (except under unusual 
conditions) . 

Many facts about gases are readily interpreted by this 
theory. For example, compressibility, diffusion, and the 
ability to mingle with other gases find explanation in the 
conception of the constant and rapid motion, perfect elas- 
ticity, and relatively great separation of the molecules. 
Boyle's law finds explanation in the conception of the pres- 
sure produced by the incessant impacts of the independent 
molecules moving in straight lines, the varying pressure 
being due to the varying number of impacts in a given time 
produced in the total space within which the gas is confined 
— the less the space, the greater the pressure. Charles' 
law is explained by the conception of the change in the veloc- 
ity of the molecules due to a change in temperature, the 
uniformly varying volume (at a constant pressure) being due 
to the varying frequency (and consequently varying energy 7 ) 
of the impacts of the molecules — the more frequent the 
number of impacts, the greater the pressure in a constant 
volume, or what is the same thing, the greater the volume 
at a constant pressure. The law of Gay-Lussac (that is, 
volumes of combining gases are in a simple integral relation 
to each other and to the total volume of the gaseous product) 
must be interpreted jointly by the atomic theory and the 
kinetic theory. According to the atomic theory chemical 
union occurs between atoms, while according to the kinetic 
theory gases consist of molecules. Now, when gases com- 
bine, is the combination between atoms or molecules? Let 
us attempt to answer this question by considering the forma- 
tion of water by the combination of oxygen and hydrogen 
gases. According to Gay-Lussac's law two volumes of hy- 
drogen and one volume of oxygen produce two volumes of 



234 INORGANIC CHEMISTRY 

water vapor. The simplest assumption (and the one actually 
made by Dalton when this question was first asked) is that 
equal volumes of elementary gases contain the same number 
of atoms. Let the number in a unit volume be X. Then 
in our illustrative case 2 X atoms of hydrogen unite with X 
atoms of oxygen to form 2 X particles of water vapor ; i.e. 
each particle of water vapor contains half an atom of oxygen ! 
But according to the atomic theory there are no fractions 
of atoms, therefore we must abandon this incorrect assump- 
tion that equal volumes of gases contain an equal number of 
atoms. Another assumption can be made, viz. that equal 
volumes of gases contain the same number of molecules. 
This assumption was first made by Avogadro, an Italian 
physicist, and is still known and accepted in a slightly modi- 
fied form as Avogadro's hypothesis (see below). If this as- 
sumption is made, then the facts summarized by Gay-Lussac's 
law can be satisfactorily explained. Let the number of 
molecules in a given volume of a gas be Z. Applying our 
assumption to the previous illustration, 2 Z molecules of 
hydrogen combine with Z molecules of oxygen to form 2 Z 
molecules of water vapor. Now a molecule of hydrogen and 
of oxygen each consists of at least two atoms (as will soon be 
shown). Therefore the illustration can be interpreted as 
follows : The given volume contains Z molecules. The 2 
volumes of hydrogen contain 2 Z molecules or 4 Z atoms, 
while the 1 volume of oxygen contains Z molecules or 2 Z 
atoms ; or 6 Z atoms all together. When these volumes 
unite, 2 volumes of water vapor are formed which contain 2 Z 
molecules of water vapor. Now if we assume (see below) 
the fact that a molecule of water vapor consists of 2 atoms 
of hydrogen and 1 of oxygen, it is obvious that not only are 
the requisite number of atoms available to form 2 molecules 
of water vapor, but that exactly 2 molecules must be formed. 
In a few words, if we accept the view that gaseous reactions 



ATOMIC AND MOLECULAR WEIGHTS — VALENCE 235 

take place between molecules, then Avogadro's hypothesis 
follows as a logical conclusion. 

Avogadro's hypothesis is usually stated thus : — 
There is the same number of molecules in equal volumes of 
all gases at a given temperature and pressure. 

This hypothesis means, for example, that a liter of hydrogen 
and a liter of oxygen at the same temperature and pressure 
contain the same number of molecules. It therefore fol- 
lows that the weights of the molecules in equal volumes of 
these two gases (and all other gases) at a given temperature 
and pressure are in the same ratio as the weights of their 
equal volumes. Hence to find the relative weights of gaseous 
molecules, it is only necessary to determine and compare the 
actual weights of equal volumes of the gases at a given tem- 
perature and pressure. Furthermore, if we express the 
weights of molecules in terms of the standard adopted for 
atomic weights, we have found the molecular weight of 
the substance. It is clear, then, that the molecular weights 
of elements and compounds in the gaseous state can be 
found by (1) assuming the kinetic theory and Avogadro's 
hypothesis, and (2) determining experimentally the rela- 
tive weights of gases. 

The hypothesis of Avogadro was proposed in 1811. But it was 
not favorably received nor was it utilized in finding molecular 
weights until about 1858. At this time it was shown by Canniz- 
zaro, a countryman of Avogadro, to be a reliable hypothesis, since 
its application yielded molecular weights in almost complete agree- 
ment with the weights found by other methods. Or more strictly, 
the atomic weights of the elements derived by all methods (includ- 
ing the method which involved Avogadro's hypothesis) were uni- 
form, accurate, and on a consistent theoretical basis. Although this 
hypothesis cannot be verified by methods usually used in chemical 
investigations and does not have the certainty of a demonstrable 
law, it is in harmony with the laws of gases and can be logically 
deduced from the kinetic theory of gases. 



236 INORGANIC CHEMISTRY 

Determination of Molecular Weights by the Vapor Density 
Method. — The density of a gas or vapor is its relative weight, 
i.e. its weight in terms of a standard. Thus, if the density of 
hydrogen is accepted as 1, the density of oxygen is about 16, 
because a given volume of oxygen weighs about sixteen times 
as much as an equal volume of hydrogen at the same tem- 
perature and pressure. Similarly, if the density of air is 
accepted as 1, the density of oxygen is about 1.1 because the 
weights of equal volumes under the same conditions are in 
this ratio. We may, however, choose any standard, such as 
hydrogen = 2 or oxygen = 32, since the different values can 
be readily transformed into each other when we know the 
numerical relation of the standards. 

For many years hydrogen was the standard gas for ex- 
pressing vapor density, and the molecular weights of gases 
and vapors were found by multiplying their vapor densities 
by 2, thus : — 

Molecular Weight = Vapor Density referred to Hydrogen X 2. 

The vapor density is multiplied by 2 because the molecular 
weight of hydrogen is 2; and the molecular weight of hydro- 
gen is 2 because a molecule of hydrogen contains at least two 
atoms each having the atomic weight 1. The conclusion 
that the hydrogen molecule contains at least two atoms is 
based on the following : — 

One volume of hydrogen combines with one volume of 
chlorine to form two volumes of hydrogen chloride (HC1). 
Suppose the volume of hydrogen contains 1000 molecules. 
Then, according to Avogadro's hypothesis, the equal volume 
of chlorine will contain 1000 molecules, while the two volumes 
of the product will contain 2000 molecules of hydrogen chlo- 
ride. That is : — 

1000 molecules of Hydrogen + 1000 molecules of Chlorine 

= 2000 molecules of Hydrogen Chloride. 



ATOMIC AND MOLECULAR WEIGHTS — VALENCE 237 

Now since every molecule of hydrogen chloride contains at 
least one atom each of hydrogen and chlorine, the 2000 
molecules must contain at least 2000 atoms each of 
hydrogen and chlorine. But the 2000 atoms of hydrogen 
and of chlorine were provided by the 1000 molecules of 
hydrogen and the 1000 molecules of chlorine. Therefore 
each molecule of hydrogen and of chlorine must contain at 
least two atoms. There is no evidence which leads us to 
believe that the chlorine or the hydrogen molecule contains 
more than two atoms. That is, no case is known in which a 
given volume of hydrogen gas furnishes the material for more 
than two volumes of the gaseous product ; the same is true 
of chlorine. On this ground we base our belief that every 
hydrogen and every chlorine molecule contains only two 
atoms. 

There are good reasons, however, for adopting oxygen 
gas = 32 as the standard for expressing vapor density. The 
atomic weight of oxygen is 16, and this is the international 
standard for atomic weights. Hence if we express molecular 
weights in terms of the atomic weight of oxygen, these values 
then become a consistent part of the international system of 
expressing the quantitative aspects of chemical change. 
The molecular weight of oxygen is 32, because a molecule of 
oxygen contains two atoms (each weighing 16). This con- 
clusion is based on the following : — 

Two volumes of hydrogen unite with one volume of oxygen 
to form two volumes of water vapor. Suppose a single vol- 
ume contains 1000 molecules. Then the two volumes of 
water vapor must contain 2000 molecules and each mole- 
cule must contain at least one atom of oxygen, or 2000 
atoms of oxygen in all. Since 2000 atoms of oxygen were 
furnished by the 1000 molecules of oxygen, each molecule 
of oxygen must contain at least two atoms; and no facts 
lead us to believe that there are more than two atoms. 



238 INORGANIC CHEMISTRY 

As the atomic weight of oxygen is 4$, the molecular weight 
is 32. 

A method of determining the molecular weight of a gas 
(or a volatile substance) is now clear, i.e. multiply the vapor 
density referred to oxygen by 32. 

Another method is used, known as the gram-molecular 
volume method. We have already seen (page 138) that 
the gram-molecular weight of a substance is the number of 
grams numerically equal to the molecular weight. (The 
terms molar weight, formula weight, and mole are sometimes 
used instead of gram-molecular weight.) E.g. the gram- 
molecular weight of oxygen is 32 gm. Now since 1 1. of 
oxygen weighs 1.429 gm., 22.4 1. approximately (32 -f- 1.429 
= 22.393) is the volume of oxygen occupied by 32 gm. 
According to Avogadro's hypothesis equal volumes of all 
gases at the same temperature and pressure contain the 
same number of molecules. Hence 22.4 1. of any gas will 
have a weight which will not only show how many times 
heavier the gas is than oxygen but which will also be the 
molecular weight of the gas referred to oxygen as 32. This 
volume (22.4 1. at 0° C. and 760 mm.) of a gas is called its 
gram-molecular volume. Hence the second method of deter- 
mining the molecular weight of a gas is to find the weight 
in grams of 22.4 1. (at 0° C. and 760 mm.) of the gas. 

Experimental Determination of Molecular Weights. — 

The weight of one liter of oxygen has been carefully deter- 
mined. Hence the simplest method of determining the 
molecular weight of a gas or vapor would appear to be 
merely to find the exact weight of a liter of the gaseous or 
vaporized substance and make the necessary calculation. 
It is more convenient experimentally, however, to find the 
volume of air displaced by the vapor of a known weight of 
a substance and then calculate the weight of 22.4 1. 



ATOMIC AND MOLECULAR WEIGHTS — VALENCE 239 




A vapor density method frequently used is one devised by Vic- 
tor Meyer. A simplified form of the apparatus is shown in Figure 
26. The bulb B of the inner tube is heated to 
a constant temperature by the vapor of the liquid 
boiling in the larger tube A. The gas measur- 
ing tube D, filled with water, is then inverted 
in the vessel of water E over the end of the 
capillary tube C. Finally, a weighed quantity 
of the substance (in a small bulb or bottle) 
is introduced into B by quickly removing and 
replacing the stopper F ; a wad of glass wool or 
asbestos at the bottom of B prevents the tube 
from being broken. The substance soon vapor- 
izes, and the vapor forces its own volume of 
air into the gas tube E. When the substance 
is completely vaporized, the volume of air in 
D is measured and reduced to the volume it 
would occupy if it were a dry gas at 0° C. 
and 760 mm. From the corrected volume and 
the weight of the substance the weight of 22.4 1. 
is calculated. For example, .1008 gm. of chloro- 
form displaced 18.93 cc. of air (corrected 
volume). If 18.93 cc. of chloroform vapor 
weigh .1008 gm., 22.4 1. will weigh 118.3 gm. 
That is, according to this experiment, the molec- \ }A 

ular weight of chloroform is 118.3 (the exact 
value being 119.5). Fig. 26.— Appara- 

tus for determin- 

The vapor density method is limited to ing vapor density * 
gaseous or volatile substances. Other 
methods are now available, viz. the freezing point and boiling 
point methods. 

In Chapter IX it was stated that the freezing point of a 
solution is lower than the freezing point of the solvent, and 
that the depression of the freezing point in dilute solutions 
is approximately proportional to the concentration of the 
solution in the case of all substances which are not ionized or 
do not unite with the solvent. Furthermore, by extending 
the results obtained by experiments with dilute solutions, it 



240 



INORGANIC CHEMISTRY 



has been found that if a mole (i.e. a number of grams numer- 
ically equal to the molecular weight) of substances which 
depress the freezing point normally is dissolved in 1000 gm. 
of water, the freezing point of all such solutions is depressed 
the same number of degrees ; i.e. each solution freezes at 
approximately — 1.86° C. This means, for example, that a 
solution of 342 gm. of sugar (C12H22O11) in 
1000 gm. of water freezes at approxi- 
mately — 1.86° C. Since this number 
( — 1.86) is the same for all solutions 
containing the molecular weight in grams 
in 1000 gm. of water, it is sometimes 
called the molecular depression constant 
(K). Now, if we find by experiment the 
amount of depression of the freezing 
point caused by a solution of known 
concentration, the molecular weight of 
the solute is readily calculated. An 
example will make this point clear. It 
was found that a solution of 50 gm. of 
methyl alcohol in 1000 gm. of water froze 
at approximately— 2.90° C. A solution 
containing one mole (i.e. a gram-molec- 
ular weight) of methyl alcohol in 1000 
gm. of water would freeze at approxi- 
mately — 1.86° C. Hence the molecular 
weight in grams is found by the pro- 
portion — 

50:z::-2.90: -1.86 

or x = 32. That is, 32 is the molecular 
weight of methyl alcohol. 







Fig. 27. — Beckmann 
apparatus for finding 
molecular weights by 
depression of the 
freezing point. 



This method of finding molecular weights was first applied ex- 
tensively by Raoult (1830-1901) and is sometimes called the cryo- 



ATOMIC AND MOLECULAR, WEIGHTS — VALENCE 241 

scopic method. The experiment may be performed in an apparatus 
devised by Beckmann (Fig. 27). The solvent is placed in the inner 
tube A provided with a side tube E through which the solute is 
introduced. A very accurate and sensitive thermometer T passes 
through a stopper into the inner tube, which is also provided with 
a stirrer S. An outer tube C serves as an air jacket, and is sur- 
rounded by the freezing mixture, which is placed in the large vessel 
B. A weighed quantity of the solvent is put into the tube A, and 
its freezing point is carefully determined. Then a weighed quan- 
tity of the solute is introduced through the side tube E, and the 
freezing point of the solution is determined. The difference be- 
tween the two freezing points is the depression, and from this de- 
pression the molecular weight can be calculated. Let us take an 
example. Suppose 4.98 gm. of sugar (C12H22O11) are dissolved in 
96.9 gm. of water and the depression is .287° C. Since 96.9 gm. of 
water contain 4.98 gm. of sugar, 1000 gm. of water would contain 
51.39 gm. of sugar. If 51.39 gm. of sugar cause a depression of 
.287°, then the number of grams which would cause the molecular 
depression can be found by the proportion — 

51.39: .287 ::x: 1.86 

or x = 333. The correct value is 342. 

Molecular weights can also be determined by an analogous 
method known as the boiling-point method. 

Exact and Approximate Molecular Weights. — The mo- 
lecular weights found by experiment are only approximate ; 
that is they are not exactly, though often very nearly, equal 
to the sum of the exact atomic weights in one mole- 
cule. The difference is due partly to the difficulties in 
making accurate measurements of temperature, partly to im- 
purities in the substances and slight defects in the experi- 
mental method, and partly also to the erroneous assumption 
that Avogadro's hypothesis is absolutely true for all tempera- 
tures. But these errors, however, do not affect the validity 
of the general result. Most molecular weights used in chemi- 
cal discussions and calculations are calculated molecular 
weights. That is, after the composition of the substance has 



242 INORGANIC CHEMISTRY 

been determined by analysis, the molecular weight found by 
experiment is slightly changed so that it will equal the sum 
of the weights of the atoms in a single molecule. It should 
be noted that the failure to find exact molecular weights by 
experiment is not a serious misfortune. The goal is atomic 
weights, and if these are accurately determined (as they can 
be), the exact molecular weight is readily found by merely 
adding the atomic weights corresponding to the number of 
atoms in the molecule. But, as repeatedly stated, approxi- 
mate molecular weights must be known in order to find the 
atomic weights of the elements. It should be noted further 
that the methods of finding molecular weights described in 
the foregoing pages apply only to substances in the gaseous 
or dissolved state. No experimental method is known for 
determining the molecular weights of substances in the 
solid state. Therefore, unless there is evidence to the con- 
trary, the molecular weight of a substance in the solid state 
is assumed to have the same value as that found by the usual 
methods. 

Relation of Atomic Weights to Molecular Weights and 
Equivalent Weights and Determination of Atomic Weights. 

— It has been shown that the atomic weights of the elements 
bear a simple numerical relation to molecular weights and 
equivalent weights. Furthermore, the methods of deter- 
mining and calculating both molecular and equivalent weights 
have been described. It is now appropriate to discuss the 
methods of selecting that weight known as the atomic 
weight, which bears the correct relation to the equivalent 
weight on the one* hand and the molecular weight on the 
other. 

The first method we shall consider might be called the 
minimum weight method. It consists in (a) determining by 
appropriate chemical and physical methods the molecular 



ATOMIC AND MOLECULAR WEIGHTS — VALENCE 243 

weights of several representative compounds of the element, 
then (b) finding the weight of the element contained in the 
molecular weight of each compound, and finally (c) select- 
ing the minimum value from the values obtained in (b). 
When this process is applied to several of the well known 
elements, the results may be tabulated (in round numbers) 
as follows : — 



Compound 


82 
■ S 


z 

© 

M 

O 


© 
o 

pq 


z 

5 


z 

© 

o 

H 


z. 
o 

M 

< 


i-i 


Water 


18 


16 


2 














Hydrogen Dioxide 




34 


32 


2 


— 


— 





— 


Hydrogen Chloride 




36.5 


— , 


1 


35.5 


— 





— 


Ammonia . . . 




17 


— 


3 


— 


14 





— 


Nitric Acid . . 




63 


48 


1 


— 


14 





— 


Nitrous Oxide . . 




44 


16 


— 


— 


28 





— 


Nitric Oxide . . 




30 


16 


' — 


— 


14 





— 


Nitrogen Dioxide . 




46 


32 


— 


— 


14 





— 


Carbon Monoxide 




28 


16 


— 


— 


— 


12 


— 


Carbon Dioxide . 




44 


32 


— 


— 


— 


12 


— 


Methane ... 




16 


— 


4 


— 


— 


12 


— 


Ethylene . . . 




28 


— 


4 


— 


— 


24 


— 


Acetylene . . . 




26 


— 


2 


— 


— 


24 


— 


Ether 




74 


16 


10 


— 


— 


48 


— 


Ethyl Alcohol . . 




46 


16 


6 


— 


— 


24 


— 


Chloroform . . . 




119.5 


— 


1 


106.5 


— 


12 


— 


Carbon Tetrachlorid* 




154 


— 


— 


142 


— 


12 


— 


Cyanogen Chloride 




61.5 


— 


— 


35.5 


14 


12 


* — 


Sulphur Dioxide . 




64 


32 


— 


— 


— 


— 


32 


Sulphur Trioxide . 




80 


48 


— 


— 


— 


— 


32 


Carbon Disulphide 




72 


— 


— 


— 


— 


12 


64 


Hydrogen Sulphide . 




34 


— 


2 


— 


— 


— . 


32 


Sulphuryl Chloride 




135 


32 


— 


71 


— 


— 


64 


Minimum weight of each element 


16 


1 


35.5 


14 


12 


~32~ 



244 INORGANIC CHEMISTRY 

In this table columns one and two contain the names of 
the compounds and their approximate molecular weights ; 
the other columns contain the parts of the molecular weights 
that belong to the atoms of the elements in a molecule of 
the compound. For example, the procedure in the case of 
water is as follows : (1) By experiment the approximate 
molecular weight of water vapor is found to be 18 ; (2) by 
analysis the compound is shown to contain 88.82 per cent 
oxygen and 11.18 per cent hydrogen; and (3) the products 
of 18 and these percentages are 16 and 2 respectively. An 
examination of the weights of the elements shows that in 
each column (1) the minimum weight is = 16, H = 1, 
CI = 35.5, N = 14, C = 12, and S = 32, and (2) the other 
weights are simple multiples. These facts are significant. 
The smallest weights must be the weights of single atoms, 
for it is highly probable that one or more compounds in a 
representative group will contain only one atom of a given 
element ; and in these compounds, of course, the part of the 
molecular weight apportioned to the element in question is 
the atomic weight. In the other compounds that contain 
this element the part of the molecular weight apportioned 
to the element will be a multiple of the atomic weight ; 
obviously these compounds contain two or more atoms of 
the element. Thus, in hydrogen dioxide the weight appor- 
tioned to oxygen is twice that in water, and hence a mole- 
cule of hydrogen dioxide contains two atoms of oxygen. 

A comparison of the atomic weights and the equivalent 
weights of the elements shows that the atomic weight of 
an element is equal to the equivalent weight or to a small 
integral multiple of it. (See table on p. 231.) This rela- 
tion now demands a fuller explanation than hitherto given. 
The equivalent weight, as previously stated, is a number 
obtained by expressing in a special way the per cent in 
which elements combine or exchange places. It is simply 






ATOMIC AND MOLECULAR WEIGHTS — VALENCE 2i5 

the number of parts by weight of an element which com- 
bine with or displace 8 parts by weight of oxygen. But 
the equivalent weight is not the unit used in chemistry to 
express the composition of compounds or the quantitative 
relations of the elements. The atomic weight is the quan- 
titative unit. Hence the equivalent weights must be multi- 
plied by certain integers in order to transform the equivalent 
weight of an element into the corresponding atomic weight. 
That is to say, the equivalent weights are the fundamental 
empirical proportions, while the atomic weights are the ad- 
justed gravimetric units deduced from the equivalent weights 
on the one hand and from the molecular weights on the 
other hand. 

Another method for determining atomic weights, appli- 
cable to the solid elements only, is known as Dulong and 
Petit's method because it utilizes an approximate law an- 
nounced and applied by them about 1819. The law, com- 
monly called the law of specific heats, may be stated as 
follows : — 

The product of the specific heat and atomic weight of the solid 
elements is approximately equal to 6.25. 

By specific heat is meant the quantity of heat necessary to 
raise the temperature of a substance one degree compared 
with the quantity necessary to raise the temperature of the 
same w T eight of w r ater one degree. If the same quantity of 
heat is imparted to equal weights of water and mercury, 
the temperature of the mercury will be much higher — about 
32 times higher than that of the water. That is, the mercury 
requires only about -^ as much heat as the water. In other 
words, the specific heat of mercury is -^, or .0312. The 
specific heat of elements in the solid state can be found 
readily by experimental methods. The number found by 
multiplying the specific heat of a solid element by its atomic 
weight varies somewhat, but in many cases it is between 



246 



INORGANIC CHEMISTRY 



6 and 6.5 (approximately 6.25). This relation is illustrated 
by the following : — 

Table of Specific Heats 



Element 


Specific Heat 


Approximate 
Atomic Weight 


Product 


Calcium 

Copper 

Iron 

Lead 

Potassium 

Sodium 

Sulphur 

Tin 

Zinc 


.170 
.095 
.114 
.031 
.166 
.293 
.178 
.055 
.094 


40 

63.5 

56 
207 

39 

23 

32 
119 

65 


6.8 

6.03 

6.38 

6.41 

6.47 

6.73 

5.7 

6>54 

6.11 



Dulong and Petit's law is only an approximation, and it 
serves merely to check values obtained by other methods 
and to show whether the atomic weight of an element is a 
multiple of the equivalent, or identical with it. The use of 
the law in checking atomic weights may be illustrated as 
follows: The specific heat of silver was found to be .057; 
if 6.25 is divided by this number, the quotient is about 109. 
This result shows that the atomic weight of silver is certainly 
not 55 or 218, but is approximately 109 (the. exact value 
being 107.88). Again, the atomic weight of uranium may be 
238 or 119 according to the chemical analysis; but only the 
former conforms to Dulong and Petit's law, and hence it is 
accepted as the approximate atomic weight. 

When the approximate atomic weight of an element has 
been chosen on the basis of the foregoing principles, the task 
still remains to determine the accurate value of this impor- 
tant weight by chemical analysis. The general method em- 
ployed can be illustrated by a determination made by the 
American chemist Richards, whose work is very accurate. 



ATOMIC AND MOLECULAR WEIGHTS — VALENCE 247 

He found that 28.26299 gm. of silver chloride were formed 

from 21.27143 gm. of silver. He accepted AgCl as the 

formula of silver chloride and 107.880 as the atomic weight 

of silver, and calculated the atomic weight of chlorine 

thus : — 

28.26299 - 21.27143 = 6.99156 

Wt. of Silver : Wt. of chlorine : : At. wt. of silver : At. wt. of chlorine 
21.27143 : 6.99156 :: 107.880 x 

x = 35.458 

The international atomic weight of chlorine (35.46) is 
based on this and other determinations made by the same 
chemist. 

The exact determination of atomic weights is a difficult 
task. And although the utmost care is used in purifying 
the chemicals and performing the analysis, the results of 
different experimenters do not always exactly agree. There- 
fore an international committee was chosen several years 
ago to select the most accurate atomic weights of the ele- 
ments. These weights are embodied in a table published 
annually and called the International Table of Atomic 
Weights. The entire table is given in the Appendix, § 5, and 
a supplementary table is given on the inside of the back 
cover. In the latter table the accepted atomic weights 
are placed in one column and the approximate, values in 
another. The approximate atomic weights are often suffi- 
ciently accurate for general reference and in making chemical 
calculations; they may be used in solving the problems in 
this book. 

The methods and principles used in determining the atomic weight 
of an element can be reviewed by the general plan which would be 
followed out in the case of zinc (if its atomic weight were unknown), 
(a) The equivalent of zinc is found by experiment to be 32.68. 
Therefore the atomic weight is 32.68, or some multiple of this num- 



248 INORGANIC CHEMISTRY 

ber. (b) The molecular weight of zinc chloride is found by its vapor 
density to be 136. (c) Analysis of zinc chloride yields 47.8 per cent 
of zinc. Therefore, 47.8 per cent of 136, or 65.08, is zinc. That is, 
this number 65.08 is the weight of the smallest part of zinc in zinc 
chloride, and it may be the weight of one atom, (d) The specific 
heat of zinc is about .094. Applying Dulong and Petit's law, the 
approximate atomic weight is found to be 66.4 (i.e. 6.25 -*• .094). 
Therefore, the atomic weight is about 65, and not 32.68. (e) Careful 
analysis of zinc compounds shows that the atomic weight of zinc 
(on the basis O == 16) is 65.37. 

Determination of Formulas of Compounds. — The formula 
of a compound is an expression of its composition; that is ; 
it is a group of symbols which not only expresses the pro- 
portions of the weights of the elements in a compound, but 
also the number of atoms whose sum equals the molecular 
weight. In a word, a formula is a molecular formula, pro- 
vided, of course, the molecular weight is known. For ex- 
ample, the proportion of hydrogen to chlorine in hydrogen 
chloride is 1 to 35.5, and the molecular weight is known to 
be 36.5; therefore, there must be one atom each of hydrogen 
and chlorine in a molecule, and the formula is HC1. 

Again, the proportion of carbon to hydrogen in a certain 
hydrocarbon is 12 to 2, and its vapor density (referred to 
oxygen) is 5.09. The formula .would be CH 2 according to the 
proportions by weight. But the molecular weight of such a 
compound would be only 14 (i.e. 12 + 2), whereas the vapor 
density of the original hydrocarbon necessitates the molec- 
ular weight 162.88 (i.e. 32 x 5.09), which is nearly twelve 
times the weight corresponding to CH 2 . Hence the molec- 
ular formula is not CH 2 , but Ci 2 H 2 4. Again, the proportion 
of hydrogen to oxygen in hydrogen peroxide is 1 to 16, and 
the formula HO would express this proportion. But the 
molecular weight of hydrogen peroxide has been found by 
the freezing point method to be nearly 34. Therefore, 
H 2 2 , not HO, must be the formula, because H 2 2 corre- 



ATOMIC AND MOLECULAR WEIGHTS — VALENCE 249 

sponds to the molecular weight as well as to the proportions 
by weight. 

If the composition of a compound is known, the smallest 
number of atoms corresponding to the composition can be 
readily calculated by dividing the per cent of each element 
in the compound by the atomic weight (and, if necessary, 
reducing the quotients to the smallest w T hole numbers) . That 
is, we merely transform the percentage into numbers which 
express the atomic relations by distributing the propor- 
tional parts among the elements according to that method 
of expressing composition adopted in chemistry. Some 
examples will make this point clear. According to analysis, 
a compound contains 40 per cent of calcium, 12 per cent of 
carbon, and 48 per cent of oxygen. Dividing each per cent 
by the proper atomic weight, we have : — 

40 -^ 40 = 1 
12-5-12 = 1 

48-^16=3 

That is, one molecule of this compound contains (at least) 
one atom each of calcium and carbon, and three atoms of 
oxygen. Hence, the simplest formula is CaC0 3 . Again, a 
compound was found to contain 2.04 per cent of hydrogen, 
32.65 per cent of sulphur, and 65.31 per cent of oxygen. 
Proceeding as in the previous case, we have : — 

2.04-1 =2.04 
32.65-5-32 = 1.02 

65.31-^16 = 4.08 

Reducing these quotients to the smallest set of integral 
numbers, we have : — 

2.04-5-1.02 = 2 

1.02-*- 1.02 = 1 

4.08-5-1.02 = 4 



250 INORGANIC CHEMISTRY 

That is, one molecule of this compound contains two atoms 
of hydrogen, one of sulphur, and four of oxygen. Hence 
the simplest formula is H 2 S0 4 . If the molecular weight of 
a compound cannot be found, then the simplest formula 
(found as in the above cases) is usually accepted as the 
molecular formula. Finally, a compound was found by 
analysis to contain 92.3 per cent of carbon and 7.7 per cent 
of hydrogen; the vapor density (referred to oxygen) was 
2.4375. Proceeding as above, we have : — 

92.3-^12 = 7.7 or 1 

7.7 h- 1 = 7.7 or 1 
That is, the compound contains at least one atom of car- 
bon and hydrogen, and would have the formula CH, if 
nothing were known about its molecular weight. The vapor 
density 2.4375 requires the molecular weight 78, which is six 
times the weight (13) corresponding to the formula CH. 
Hence the correct formula of this compound is not CH 
but C 6 H 6 . 

To recapitulate : The simplest formula of a compound 
is found by dividing the per cent of each element by its 
atomic weight and reducing these quotients to the smallest 
whole numbers (if necessary) ; if the molecular weight is not 
known, this formula is accepted as the molecular formula. 
The molecular formula of a compound is found by three 
steps : (a) Find the simplest formula, (b) divide the molecular 
weight by the sum of the weights of the atoms in the simplest 
formula, (c) multiply the integral numbers of the simplest 
formula by the quotient obtained in (b). (Compare illustra- 
tion in preceding paragraph.) 

Molecular Weights and Molecular Formulas of Elements. 

— Several elements are gases at ordinary temperatures and 
others can be vaporized by heating. Hence their molecular 
weights can be determined by the vapor density method. 



ATOMIC AXD MOLECULAR WEIGHTS — VALENCE 251 

When this is done, the results indicate that the elements 
fall into three classes. (1) The gaseous elements already 
studied, as well as some others, have molecular weights 
which are twice the atomic weight ; that is, the molecule 
consists of two atoms, and their molecular formulas are, for 
example, 2 , H 2 , Cl 2 , N 2 . (2) The molecular weights of several 
metallic elements and certain gaseous elements are identi- 
cal with their atomic weights ; that is, the molecule consists 
of one atom, and the molecular formula is the same as the 
atomic symbol, e.g. Na, K, Zn, Hg, Cd (cadmium), A (argon), 
He (helium), and Ne (neon). (3) The molecular weights of 
certain elements decrease with rise of temperature ; e.g. at 
lower temperatures, molecules of iodine, sulphur, and phos- 
phorus are represented by I 2 , S 8 , and P 4 , and at higher tem- 
peratures by I, S 2 , and P 2 . At intermediate temperatures 
the two kinds of molecules are in equilibrium. 

Molecular Equations. — Since reactions between gases are 
between molecules, equations representing such reactions 
should be written in the molecular form. For example, 
since a molecule of hydrogen has the formula H 2 and of 
oxygen the formula 2 , the molecular equation for the forma- 
tion of water vapor from hydrogen and oxygen is — 

2 H 2 + 2 = 2 H 2 

It is read thus : Two molecules of hydrogen unite with one 
molecule of oxygen to form two molecules of water vapor. 
Since most elementary gases consist of molecules, such 
equations correctly represent the actual substances in- 
volved. It should be noted, however, that the proportions 
by weight are the same as in the simpler or atomic form 
of the equation. Molecular equations are sometimes called 
volume equations or gas equations. Thus, the equation 

H 2 + Cl 2 = 2 HC1 



252 INORGANIC CHEMISTRY 

means that one volume each of hydrogen and chlorine unite 
to form two volumes of hydrogen chloride, or more fully 



H 2 


+ 


Cl 2 


2HC1 


2 gm. 




71 gm. 


73 gm. 


22.4 1. 




22.4 1. 


44.8 1. 


1 volume 




1 volume 


2 volumes 


1 molecule 




1 molecule 


2 molecules 



The equation may be written in the latter form because 
equal numbers of molecules represent equal volumes. It is 
very convenient to remember that in molecular or gas equa- 
tions one molecule represents one volume of a gas. 

Valence 

A classification of the formulas of chemical compounds 
shows certain regularities. Instead of a vast number of 
unrelated formulas, there are really only a few groups, 
especially if attention is confined to inorganic compounds. 
Take, for instance, some binary compounds of hydrogen. 
They fall into the following classes : — 

HC1 H 2 NH 3 CH 4 

HBr H 2 S PH 3 SiH 4 

HI AsH 3 

HF . SbH 3 

Again, if some oxides are considered, we have the follow- 
ing : — 



Na 2 


CuO 


A1 2 3 


Mn0 2 


P 2 O s 


SO s 


K 2 


MgO 


P 2 3 


S0 2 


As 2 5 


CrO s 


Ag 2 


CaO 


As 2 3 


C0 2 







We might also group certain acids and salts thus : — 
HC1 H 2 S0 4 HNO3 H 3 P0 4 

NaCl Na 2 S0 4 NaN0 3 Na 3 P0 4 

CaCl 2 CaS0 4 Ca(N0 3 ) 2 Ca 3 (P0 4 ) 2 

A1C1 3 A1 2 (S0 4 ) 3 Bi(N0 3 ) 3 Mg 3 (P0 4 ) 2 



ATOMIC AND MOLECULAR WEIGHTS — VALENCE 253 

Similarly, there are the following groups of bases : — 

NaOH Ca(OH) 2 Al(OH) 8 

KOH Ba(OH) 2 Bi(OH) 3 

If these lists were to be extended, similar regularities 
would be revealed. A careful examination of these formu- 
las leads to two conclusions. (1) Atoms of elements differ 
in the number of atoms or atomic groups of the other ele- 
ments with which they combine. Thus, in the first list 
one atom each of chlorine, bromine, iodine, and fluorine 
combines with one atom of hydrogen; one atom of oxygen 
and of sulphur combines with two of hydrogen; and so on 
through all the lists. (2) An atom of certain elements 
unites with only one atom or atomic group of certain other 
elements, with only two atoms or two atomic groups of cer- 
tain others, etc. Thus, one atom of calcium (Ca) combines 
with one atom of oxygen, with two of chlorine, with one S0 4 - 
group, and with two N0 3 -groups. Hence w T e conclude that 
atoms of the elements have a definite and characteristic com- 
bining capacity. The number which expresses the maxi- 
mum combining capacity of an atom of an element is called 
the valence of the element. 

Determination of Valence. — The valence of an element is 
found by dividing the atomic weight by the equivalent 
weight. It will be recalled that elements combine in the 
ratio of their equivalent weights. That is, if the per cent of 
each element in a compound is restated so that these per 
cents become the number of grams which unites with or 
replaces eight grams of oxygen, the composition of the com- 
pound is expressed in terms of the equivalent weights of 
the constituent elements. For example, magnesium and 
oxygen combine in the ratio of 60 and 40 per cent respec- 
tively (in round numbers). If for the 40 per cent we substi- 



254 



INORGANIC CHEMISTRY 



tute 8, then the 60 per cent becomes 12, i.e. the equivalents 
of oxygen and magnesium respectively. We have seen, 
however, that atomic weights, not equivalent weights, are 
the weights used in chemistry for expressing composition. 
The number, therefore, which expresses the maximum com- 
bining capacity of an atom of an element is merely the 
factor by which the equivalent weight is multiplied to con- 
vert it into the atomic weight. To find the valence, it is 
only necessary to divide the atomic weight by the equivalent 
weight. In brief, 

Valences Atomic Weight ' 
Equivalent Weight 

Tables of Valence. — Pursuing this method of finding 
valence, we obtain the following table : — 

A. Table of Valence of Common Elements 



Element 


Symbol 


Valence 


Aluminium . 


Al 


III 


(Ammonium) 


(NH 4 ) 


I 


Antimony 


Sb 


III in antimonous and V in antimonic compounds 


Arsenic . . 


As 


III in arsenious and V in arsenic compounds 


Barium . . 


Ba 


II 


Bismuth . . 


Bi 


III 


Boron . . . 


B 


III 


Bromine . . 


Br 


I in hydrobromic acid (HBr) and bromides 


Cadmium . . 


Cd 


II 


Calcium . . 


Ca 


II 


Carbon , . 


C 


IV 


Chlorine . . 


CI 


I in HC1 and chlorides ; V in chlorates 


Chromium 


Cr 


III; VI in chromates and dichromates 


Cobalt . . . 


Co 


II 


Copper . . 


Cu 


I in cuprous and II in cupric compounds 


Fluorine . . 


F 


I 


Gold . . . 


Au 


I in aurous and III in auric compounds 


Hydrogen 


H 


I 



ATOMIC AND MOLECULAR WEIGHTS — VALENCE 255 



A. TabEe of Valence of Common Elements — Continued 



Element 


Symbol 


Valence 


Iodine . . . 


I 


I in hydriodic acid (HI) and iodides 


Iron . . . 


Fe 


II in ferrous and III in ferric compounds 


Lead . . . 


Pb 


II ; IV in Pb0 2 


Lithium . . 


Li 


I 


Magnesium . 


Mg 


II 


Manganese . 


Mn 


II ; IV in MnOo ; VI in inanganates ; VII in 
permanganates 


Mercury . 


Hg 


I in mercurous and II in mercuric compounds 


Nickel . . . 


Ni 


II 


Nitrogen . . 


N 


1 in N 2 ; III in N 2 3 and nitrites ; V in N 2 5 , 
HNG 3 , and nitrates 


Oxygen . . 





II 


Phosphorus . 


P 


III in P0O3 ; V in P 2 5 , H3PO4, and orthophos- 
phates 


Platinum . . 


Pt 


IV 


Potassium 


K 


I 


Silicon . . . 


Si 


IV 


Silver . . . 


Ag 


I 


Sodium . . 


Na 


I 


Strontium 


Sr 


II 


Sulphur . . 


S 


II in H 2 S and sulphides ; VI in S0 3 , H 2 S04, and 
sulphates 


Tin .... 


Sn 


II in stannous and IV in stannic compounds 


Zinc . . . 


Zn 


II 



It is customary to call valence a property possessed by 
atoms of elements. But groups of atoms of different ele- 
ments often act chemically like individual atoms of ele- 
ments {i.e. they pass as a whole from compound to com- 
pound) ; hence it is customary to assign a valence to atomic 
groups like S0 4 , N0 3 , etc. These groups are often called 
radicals, and are analogous to NH 4 and OH. Tabulating 
the common radicals and a few elements, we have : — 



256 



INORGANIC CHEMISTRY 



B. Table of Valence of Certain Elements and Radicals 



Chemical Name of Compound which contains 
Element or Kadical 



Symbol or 
Formula 



Valence 



Acetate . . . . 

Bromide . . . . 
Carbonate (normal) 
Carbonate (acid) 

Chlorate . . . . 

Chloride . . . . 

Chromate . . . . 

Cyanide . . . . 

Bichromate . . . 

Ferricyanide . . . 

Ferrocyanide . . . 

Fluoride . . . . 

Hydroxide . . . . 

Iodide . . , . . 

Manganate . . . 

Nitrate 

Nitrite 

Oxide 

Permanganate . . 

Phosphate (ortho) . 

Silicate (meta) . . 

Sulphate . . . . 

Sulphate (acid) . . 

Sulphide . . . . 

Sulphite (normal) . 

Sulphite (acid) . . 



C2H3O2 
Br 

C0 3 

HCO3 

CIO3 

CI 

Cr0 4 

CN 

Cr 2 7 

Fe(CN) 6 

Fe(CN) 6 

F 

OH 

I 

M11O4 

N0 3 

N0 2 

O 

Mn0 4 

P0 4 

Si0 3 

S0 4 

HS0 4 

s 

so 3 

HS03 



I 

I 

II 

I 

I 

I 

II 

I 

II 

III 

IV 

I 

I 

I 

II 

I 

I 

II 

I 

III 

II 

II 

I 

II 

II 

I 



It is obvious from these tables that the valence of some 
elements varies. Many elements have a fixed valence, but 
several of the common elements have a valence which is 
determined by the element with which they are united and 
the conditions under which the union occurred. The 
valence of an element is never less than 1 nor more than 8. 
The valence of most radicals is fixed. 



ATOMIC AND MOLECULAR WEIGHTS — VALENCE 257 

Valence Terms. — Elements and groups with the valence 
I are called monads or univalent elements, and those with 
the valence II are called dyads or bivalent elements; simi- 
larly, elements or groups which have the valence III, IV, V, 
VI, and VII are called respectively triads, tetrads, pentads, 
hexads, and heptads, and the corresponding terms are triva- 
lent, quadrivalent, quinquivalent, hexavalent, and heptav- 
alent. 

Valence applied to Combination and Displacement. — 
(1) Elements which have the same valence combine with each 
other atom for atom. Thus, one atom of sodium combines 
with one atom of chlorine, and one atom of magnesium com- 
bines with one atom of oxygen. (2) On the other hand, ele- 
ments having a different valence usually combine with each 
other so that the total valence of the atoms of each element 
in a molecule is equal, i.e. in many compounds the total 
combining capacity of each element is equal. This condi- 
tion is often described by calling the compound saturated, 
or as one having a balanced valence. Thus, two atoms of 
hydrogen combine with one of oxygen to form one molecule 
of water (H 2 0) ; two atoms of aluminium each having the 
valence III combine with three atoms of oxygen each hav- 
ing the valence II to form one molecule of aluminium oxide 
(A1 2 3 ) ; one atom of carbon having the valence IV com- 
bines with two atoms of sulphur each having the valence II 
to form one molecule of carbon disulphide (CS 2 ). (3) The 
above rules apply to atomic groups. Thus, one NH 4 -group 
combines with one OH-group to form one molecule of ammo- 
nium hydroxide (NH 4 OH) ; two NH 4 -groups combine with 
one S0 4 -group having the valence II to form one molecule 
of ammonium sulphate ((NH 4 ) 2 S0 4 ) ; and three atoms of 
calcium each having the valence II combine with two P0 4 - 
groups each having the valence III to form one molecule of 
calcium phosphate (Ca 3 (P0 4 ) 2 ). 



258 INORGANIC CHEMISTRY 

It is convenient to interpret valence not only from the 
standpoint of combination, but also from the standpoint of 
displacement. That is, just as atoms and groups of the 
same valence combine unit for unit, or those of different 
valence combine so that the valences of each element or 
atomic group balance, so also atoms and atomic groups dis- 
place each other — unit for unit if the valence is the same, 
or equivalently if the valence is different. For example, 
when zinc displaces hydrogen in hydrochloric acid, one atom 
of zinc having the valence II displaces two atoms of hydro- 
gen each having the valence I, and the formula of the result- 
ing zinc chloride is ZnCl 2 ; similarly, one atom of sodium having 
the valence I displaces one atom of hydrogen from water, and 
the formula of the resulting sodium hydroxide is NaOH. 

Valence and Formulas. — The chief value of valence in a 
course in inorganic chemistry is the assistance it gives in 
writing formulas. If it were necessary to remember the 
formula^of every compound, the study of chemistry would 
be tedious and almost hopeless. The application of valence 
to formula writing can be best illustrated by several ex- 
amples. Suppose we wish to write the formula of mag- 
nesium chloride. From Table A the valence of magnesium 
is found to be II, and from Table B the valence of chlorine 
in chlorides is found to be I ; therefore, it is necessary to take 

two atoms of chlorine for one of magnesium in order to 

ii ii 
balance the valence, and the formula is MgCl 2 . It is help- 
ful to write the valence as shown in the formula just given, 
though such a device may be abandoned after the valence 
of each element and group has been learned. Again, sup- 
pose we wish to write the formula of lead nitrate. As in 
the previous case, from Table A the valence of lead is found 
to be II, and from Table B the valence of the N0 3 -group in 
nitrates is found to be I ; therefore it is necessary to take 



ATOMIC AND MOLECULAR WEIGHTS — VALENCE 259 

two N0 3 -groups to balance the valence of the lead, and the 

ii ii 
formula is Pb(N0 3 ) 2 . Similarly, the formula of sodium 

ii ii 
sulphate is Na 2 S0 4 because two atoms of sodium each hav- 
ing the valence I are needed to balance the valence II of 

the S0 4 -group ; and the formula of aluminium hydroxide is 

in in 

Al(OH) 3 because three OH-groups each having the valence 

I are needed to balance the valence of one atom of aluminium 
having the valence III. Furthermore, suppose we wish to 
write the formula of a compound formed by the interaction 
of aluminium and hydrochloric acid. The valence of alumin- 
ium is found from Table A to be III, and that of hydrogen 

to be I. Therefore one atom of aluminium displaces three 

in in 
atoms of hydrogen, and the formula is A1C1 3 . Finally, the 

formulas of acids, although usually learned by continued 

observation, can readily be written by utilizing Table B. 

Acids may be regarded as compounds formed by the 

union of hydrogen with radicals and elements in Table B. 

That is, to write the formula of an acid it is only necessary 

to prefix the proper number of hydrogen atoms to the atom 

or group in the table, remembering, of course, the rules for 

i i 
naming acids. Thus, the formula of chloric acid is HC10 3 , 

because the valence of C10 3 is I, and this valence is balanced 

i i 
by one atom of hydrogen; similarly, nitrous acid is HN0 2 , 

in in ii ii 

phosphoric acid (ortho) is H 3 P0 4 , silicic acid is H 2 Si0 3 . 

The student is advised to write the formulas of many com- 
pounds until valence and its application become thoroughly 
familiar. (See Exercises at the end of this chapter.) 

Representation of Valence. — One way of representing 

the valence of elements and radicals has been given, viz. 

ii i ii 

O, CI, S0 4 . Sometimes short lines are used, e.g. H — , O = , or 



260 INORGANIC CHEMISTRY 

— O — , Al =, etc. If lines are used to indicate valence in 
compounds, a single line answers for two elements. Thus, 
the formula of water is H - O - H rather than H — — H. 

Similarly, ammonia gas is N — H and calcium hydroxide is 

\H 

/OH 
Ca . The lengths of the lines have no significance, 

\OH 
since the lines do not represent intensity of attraction be- 
tween the atoms. Such formulas are called structural or 
graphic formulas, and they often serve to show what is 
known about the arrangement of the atoms in a molecule. 

Apparent Exceptions. — The formulas of some compounds 
do not conform to the conception of valence discussed above, 
e.g. carbon monoxide (CO), calcium carbide (CaC 2 ), mag- 
netic iron oxide (Fe 3 4 ), acetylene (C 2 H 2 ), and ethylene 
(C 2 H 4 ). Their interpretation must be sought elsewhere. 

Problems and Exercises 

1. The vapor densities of certain gases (referred to oxygen) are 
as follows: (a) hydrochloric acid 1.14, (b) chlorine 2.218, (c) ammonia 
.53125, (d) nitrogen .875, (e) steam .5625. Calculate the molecular 
weight of each. 

2. Calculate the simplest formula of the compounds which have 
the indicated composition : (a) N = 82.353, H = 17.647; (6) O = 30, 
Fe (iron) =70; (c) H = 1, C = 12, K (potassium) = 39, O = 48. 

3. A liter of sulphurous oxide gas (SO2) weighs 2.8672 gm. What 
is the molecular weight of this compound? 

4. If 1500 cc. of carbon monoxide gas (CO) weigh 1.8816 gm., 
what is the molecular weight of the compound? 

5. Calculate the molecular formula of the. compounds correspond- 
ing to the following data: (a) C = 73.8, H = 8.7, N = 17.1, vapor 
density = 5.03 ; (6) C = 92.3, H = 7.7, vapor density - 2.425; (c) 
C = 39.9, H = 6.7, O = 53.4, vapor density = 1.906 (Vapor density 
in each case is referred to oxygen.) 



ATOMIC AND MOLECULAR WEIGHTS — VALENCE 261 

6. What volumes of factors and products are represented by the 
equations (a) H 2 + C1 2 = 2HC1, (6) 2H 2 + 2 = 2H 2 0, (c) 3 H 2 
+ N 2 =2 NH 3 , (d) N 2 + 2 =2 NO, (e) 2 NO + 2 = 2 N0 2 ? 

7. If 91.46 gm. of metallic silver yield 121.4993 gm. of pure 
silver chloride, calculate the atomic weight of chlorine. 

8. Suppose 4.86111 gm. of ferric oxide (Fe^Os) yield 3.39995 gm. 
of iron, what is the atomic weight of iron? 

9. Assume that 1.70563 gm. of strontium bromide (SrBr 2 ) re- 
quire 1.48707 gm. of silver to precipitate the bromine as AgBr. 
What is the atomic weight of strontium? 

10. If 19.57120 gm. of mercuric chloride (HgCl 2 ) yield 14.46032 
gm. of mercury, what is the atomic weight of mercury ? 

11. A liter of the monatomic gas helium weighs .1783 gm. (at 
0° C. and 760 mm.). What is the atomic weight of helium? 

12. (a) .8606 gm. of hydrogen combined with 68.25033 gm. 
of bromine. Calculate the atomic weight of bromine. (6) 4.58644 
gm. of calcium bromide (CaBr 2 ) require 4.95025 gm. of silver to 
precipitate the bromine as AgBr. What is the atomic weight of 
calcium ? 

13. (a) If 4.59507 gm. of phosphorus trichloride (PC1 3 ) give 
14.38118 gm. of silver chloride, what is the atomic weight of phos- 
phorus? (b) The molecular weight of carbon disulphide is 76, and 
100 parts yield. 84.21 parts of sulphur; the specific heat of sulphur 
is .178. What is the atomic weight of sulphur, and how many 
atoms of sulphur in a molecule of the carbon disulphide? 

14. If 2.1281 gm. of N 2 4 give .6480 gm. of N 2 , what is the atomic 
weight of nitrogen? 

15. A certain weight of copper oxide, when heated in a current 
of hydrogen, lost 59.789 gm. of oxygen and formed 67.282 gm. of 
water, (a) If 0=16, what is the atomic weight of hydrogen? 
(6) If H = 1, what is the atomic weight of oxygen? 

16. (a) What is the atomic weight of mercury if the specific heat 
is 0.032? (b) Of lead if the specific heat is 0.031? 

17. (a) To form a certain compound, 1 gm. of carbon and 2.666 
gm. of oxygen are needed. Its vapor density referred to oxygen is 
1.375. What is its molecular formula? (6) A compound of car- 
bon and hydrogen contains 14.29 per cent of hydrogen, and 1 1. 
weighs 1.25 gm. Calculate its molecular formula. 

18. If 1.2 gm. of a substance dissolved in 24.5 gm. of water 
lower the freezing point 1.05°, what is the molecular weight of the 
substance ? 



262 INORGANIC CHEMISTRY 

19. Write the formula of the chloride of potassium, sodium, sil- 
ver, copper (ous), copper (ic), mercury (ous), mercury (ic), iron 
(ous), iron (ic), cadmium, zinc, tin (ous), tin (ic), calcium, barium, 
magnesium, strontium, cobalt, nickel, bismuth, aluminium, carbon, 
ammonium, antimony (two), phosphorus (two). (Two means -ous 
and -ic.) 

20. Write the formula of the sulphide of K, Na, Ag, Cd, Zn, Ca, 
Ba, Co, Ni, NH 4 , As (two), Sb (two), Sn (two), Pb, Fe (ous). 

21. Write the formula of the sulphate of K, Na, Ag, Cu, Fe (ous), 
Cd, Zn, Pb, Ca, Ba, Co, Mg, Ni, Mn, Sr, Cr, Al, NH 4 . 

22. Write the formula of the nitrate as in Problem 21. 

23. Write the formula of the hydroxide of NH 4 , Al, K, Na, Ca, 
Mg, Ba, Fe (ic), Ni, Co, Cd, Zn, Cr, Bi. 

24. Write the formula of the potassium compound of the ele- 
ments and radicals given in Table B (under Valence, above). Do 
the same with Ca and Mg. 

25. Write the formula of copper acetate, lithium chlorate, man- 
ganese dioxide, ammonium fluoride, sodium silicate, potassium 
manganate, barium phosphate, zinc iodide, ammonium chromate, 
silver chromate, lead acetate, cobalt nitrite, magnesium oxide. 

26. Write the formula of aluminium oxide, auric chloride, fer- 
rous bromide, ferrous carbonate, mercurous nitrate, platinum chlo- 
ride, aluminium phosphate, calcium fluoride, potassium cyanide. 

27. Write the formula of phosphoric acid, silicic acid, sulphur- 
ous acid, nitrous acid, chromic acid, acetic acid, hydriodic acid. 

28. By the Victor Meyer method .1561 gm. of a compound 
expelled 32.1 cc. of dry air at 20° C. and 744 mm. Calculate its 
vapor density and its molecular weight. 

29. What volume of the component gases can be obtained by 
the decomposition of 6 1. of ammonia gas? 

30. What volume of oxygen is used up when 20 cc. of acetylene 
burn in air? (Equation is 2 C 2 H 2 + 5 2 = 4 C0 2 + 2 H 2 0.) 

31. (a) One volume of phosphorus vapor and six volumes of 
chlorine gas form four volumes of phosphorus trichloride vapor ; 
write the molecular equation for this chemical change, (b) How 
many liters of chlorine would be needed for 248 gm. of phosphorus ? 

32. Calculate the following equivalent weights : (a) .464 gm. of 
zinc gives 174 cc. of hydrogen at 18° C. and 763 mm. (b) .075 gm. 
of magnesium gives 75.6 cc. of hydrogen at 17° C. and 760.4 mm. 
(c) .17 gm. of aluminium gives 223.7 gm. of hydrogen at 16° C. and 
754 mm. 



CHAPTER XV 
Carbon and its Oxides — Carbides 

Occurrence of Carbon. — Uncombined carbon is found 
pure in nature as diamond and graphite ; in a more or less 
impure state it occurs as coal and similar substances, which 
are included in the term amorphous carbon. Combined with 
hydrogen and oxygen, and occasionally with nitrogen also, 
it is an essential constituent of animal and vegetable matter. 
Meat, starch, fat, sugar, wood, cotton, wool, wax, flour, 
albumen, and bone are familiar examples of the vast num- 
ber of natural substances which contain carbon. It is also 
a constituent of carbon dioxide and of the carbonates, such 
as limestone, chalk, marble, and dolomite. Illuminating 
gases, gasolene, kerosene (and other products obtained 
from petroleum), turpentine, alcohol, chloroform, ether, and 
similar liquids are compounds of carbon. Carbon is also a 
constituent of thousands of manufactured compounds, such 
as dyestuffs, medicines, and perfumes. 

Diamond is pure crystallized carbon. It is found in only 
a few places in the earth. When taken from the mine, 
diamonds are usually rough-looking stones ; some are 
crystals; some are rounded like peas, and many are irregu- 
lar; they are ground into special shapes and polished to 
bring out the luster and make them sparkle (Fig. 28). The 
most expensive diamonds are colorless and without a flaw, 
and are said to be " of first water " ; diamonds having a 
yellow tint are common, and occasionally a blue, pink, or 
green one is found ; a very impure variety is black. 

263 



264 



INORGANIC CHEMISTRY 



Diamond resists the action of most chemical reagents. 
It has the high specific gravity of 3.5, and is one of the hard- 
est substances. It is brittle, and may be shattered by a blow 
with a hammer. 





Crystal 



Fig. 28. 



Rough 
— Diamonds. 



Cut 



Diamonds have always been prized as gems on account of their 
beauty, rarity, and permanence. Besides being worn as jewels, they 
are used to cut glass, and the powder and splinters (known as bort) 
are used to grind and polish diamonds and other hard gems. The 
impure variety is called carbonado, and is set into the end of the 
" diamond drill," which is used extensively for boring artesian wells 
and drilling hard rocks. 

Diamonds were formerly found in gravel deposits in India, and 
in later years in Brazil. Since 1867, however, about 95 per cent of 
the diamonds of commerce have come from South Africa. They 
occur in a bluish volcanic rock along the Vaal River, and especially 
near Kimberley. Over ten tons of diamonds have been found in 
South Africa in the last twenty-five years ! 

The successive investigations of Lavoisier, Dumas, and Davy, ex- 
tending from 1772 to 1814, showed that diamond is carbon, for when 

pure diamond was burned in oxy- 
gen, the only produce was carbon 
dioxide. This result, which ad- 
mits of no doubt, has been verified 
by many famous investigators. 
Diamonds have been made. In 
1893 Moissan dissolved pure char- 
coal in melted iron, and cooled the molten mass in water. The 
surface was so suddenly cooled that a tremendous pressure was 
exerted by the expanding iron inside the crust. This pressure 
caused the cooling carbon to crystallize into diamond. The crystals 




Fig. 29. — Artificial diamonds (en 
larged) prepared by Moissan. 



CARBON AND ITS OXIDES — CARBIDES 265 

were very small, most of them were black, a few were white, but all 
had the properties of the diamond (Fig. 29). 

Large diamonds have a fascinating history, since most of them 
passed through many hands before becoming royal jewels. Until 
recently the largest royal cut diamond was the Orloff, which weighs 
194| carats, and is in the scepter of the Czar of Russia. 1 The Kohi- 
noor, which now weighs about 106 carats, is one of the crown jewels 
of England. The largest diamond ever found was called the Culli- 
nan. It was found in South Africa in 1905, and weighed about 
3025 carats (1.37 lb. avoir.). Stones cut from it are among the 
English crown jewels. 

Graphite is a soft, black, shiny solid, which is smooth 
and greasy to the touch. Pure graphite is entirely carbon. 
It occurs native in large quantities and in many places. 
One variety is found in abundance at Ticoncieroga, New 
York. Other localities are Mexico, Ceylon, Siberia, Ger- 
many, Austria, and Italy. Sometimes crystals and grains 
are found, but it usually occurs in flaky masses or slabs. 
Unlike diamond, graphite is a good conductor of electricity 
and is often used to coat molds in electrotyping. It is 
so soft that it blackens the fingers and leaves a black mark 
on paper when drawn across it. This property is indi- 
cated by the name graphite, which is derived from a Greek 
word (graphein) meaning to write. It resembles diamond 
in its insolubility in liquids at the ordinary temperature. 
Its specific gravity is about 2.2, being considerably lighter 
than diamond. When heated intensely in a current of 
oxygen, it is converted into carbon dioxide ; but it can be 
heated to a very high temperature in the air without under- 
going chemical change. Graphite was once supposed to 
contain lead, and is even now often incorrectly called " black 
lead" and plumbago. It is used to make stove polish and 
protective paints, as a lubricant where oil w T ould be de- 
composed by the heat, as the principal ingredient of the 

1 The international carat weighs 200 milligrams. 



266 INORGANIC CHEMISTRY 

graphite crucibles which withstand extremely high tempera- 
tures, and in making electrodes for electric furnaces and 
electrolytic apparatus. Large quantities of graphite are 
consumed in the manufacture of lead pencils. The graphite 
is ground to a fine powder, mixed with more or less clay, 
and then passed through perforated plates, from which the 
" lead " issues in tiny rods. These are dried, cut into proper 
lengths, baked, and then inserted in the wooden case. 

Molten iron and other metals dissolve carbon, and when 
the metals cool, the carbon crystallizes as graphite. 

Large quantities of graphite are now manufactured by heat- 
ing a special grade of coal or of coke in an electric furnace. 
(Fig. 20). The process is electrothermal, and yields a product 
that is exceptionally suitable for electrodes. Articles of 
almost any size and shape can be made of graphite ; more- 
over the graphite is very compact and can be further shaped 
by tools. 

Amorphous Carbon is a broad term, including all varieties 
of coal, coke, charcoal, lampblack, and gas carbon. They 
are the non-crystalline varieties of impure carbon, and differ 
mainly in purity, degree of fineness, and hardness. The 
word amorphous means literally " without form," and is 
often used to designate soft,- powdery substances, but more 
especially those which are uncrystallized. 

Coal is the term applied to several varieties of impure 
carbon. It may be regarded as the final product derived 
from vegetable matter which was subjected to heat and 
pressure through long geological periods. Ages ago the 
vegetation was exceedingly dense and luxuriant upon land 
slightly raised above the sea. In the process of time this 
vegetation decayed, accumulated, and slowly became covered 
with sand, mud, and water. Owing to the heat of the earth 
and the enormous pressure of the overlying deposits, the 



CARBON AND ITS OXIDES — CARBIDES 267 

complex vegetable matter decomposed slowly into carbon 
and compounds of hydrogen and carbon called hydrocar- 




Fig. 30. — Section of part of the earth's crust near Mauch Chunk, 
Pennsylvania, showing layers of coal. 

bons. The geological and chemical changes were repeated, 
and as a result we find in the earth layers or seams of car- 
bonaceous matter varying in thickness and composition 
(Fig. 30). These are the coal beds. Coal beds contain proofs 
of their vegetable origin, viz. impressions of vines, stems, 
leaves of plants, and similar vegetable substances (Fig. 31). 





Fig. 31. — Fossil found in Fig. 32. — Section of coal as seen through 
a coal bed. a microscope. 

A thin section of coal examined through a microscope 
reveals a distinct cellular structure characteristic of vege- 
table matter (Fig. 32). The transformation of vegetable 
matter into coal was an exceedingly slow process, espe- 



268 INORGANIC CHEMISTRY 

cially where the rocks were undisturbed. In some locali* 
ties the process was hastened by deep-seated changes 
incidental to the upheaval of the rocks during mountain 
building ; the gaseous products escaped almost com- 
pletely and left the carbon, together with a small pro- 
portion of mineral matter. Consequently we find several 
varieties of coal in the earth's crust. There are three main 
classes — anthracite, bituminous, and lignite, though com- 
mercially several subclasses are variously designated. They 

LIGNITE, CROCKETT, TEX. 
Fixed Carbon Moisture Ash Volatile Matter 

LIGNITE, GLENDIVE, MONT. 



W777//// // /////// // / // / / /// ///ZZ3 



M. A. 

SUB-BITUMINOUS, LAFAYETTE, COLO. 



m//////////////////////////////////A 



M. A. 

SUB-BITUMINOUS, GALLUP. N.MEX. 



j m$m///////////////////////////////// i 



F.C. M. A. 

BITUMINOUS, CARTERVILLE. ILL. 



■■■■■■iiimiiw—— ■■■in i 

F.C. M. A. V. M. 

SEMI-BITUMINOUS, POCAHONTAS, W. VA. 



F.C. M. A. V. M. 

SEMI- ANTHRACITE, SPADRA, ARK. 



F.C. 
ANTHRACITE,- PENNSYLVANIA 



F.C. M. A.V.M 

Fig. 33. — Diagram showing the progressive change in composition of 
coal. (From National Geographic Magazine.) 

differ in composition. That is, they contain different pro- 
portions of moisture, ash or mineral matter, volatile matter 
or compounds of carbon and hydrogen driven off by careful 
heating, and fixed carbon or carbon left after the removal 
of the volatile matter and ash. The progressive change in 
composition from a very poor lignite to the best anthracite 
is shown in Figure 33. The diagram also shows the com- 



CARBON AND ITS OXIDES — CARBIDES 269 

position as determined by proximate analysis, i.e. a chemical 
analysis designed to find the per cent of the components 
mentioned above and not the actual per cent of carbon, 
hydrogen, etc. From the diagram it is seen that the fixed 
carbon (F.C.) varies from about 20 per cent in the lignite 
to about 90 per cent in the anthracite, while the volatile 
matter (V.M.) varies inversely as the fixed carbon, being 
greatest in the lignite and least in the anthracite. The per 
cent of moisture (M.) diminishes from the lignite to the 
anthracite. The per cent of ash (A.) is variable. 

The properties and uses of coal differ with the composi- 
tion. Anthracite coal is hard and lustrous. It ignites 
with difficulty, burns with a slight or no flame, and pro- 
duces an intense heat. It is used mainly for domestic pur- 
poses, — heating and cooking, — especially in the eastern 
United States. Bituminous or soft coal burns with a smoky 
flame, and in burning produces considerable volatile matter ; 
some varieties form coke when heated. It is used to make 
illuminating gas, coke, and as a fuel for generating steam. 
Lignite or brown coal is the least valuable as fuel. The use 
of coal as fuel depends, of course, on the fact that consider- 
able heat is liberated when it burns; that is, it has calorific 
value. The following table shows the calorific value of 
well-known commercial grades of the three classes of coal 
in calories per gram, i.e. the number of calories liberated 
when 1 gm. is burned freely: — 

Calorific Yalue of Coal 



Class 


Calories per Gram 


Anthracite 


7724 


Bituminous 


8768 


Lignite 


4530 



270 



INORGANIC CHEMISTRY 



Coal is widely distributed in the crust of the earth, but the deposits 
vary in extent and quality. It underlies about one sixth of the area 
of the United States, the anthracite variety covering less than five 
hundred square miles in eastern Pennsylvania (Fig. 34). The United 




Fig. 34. — Map showing coal areas in the United States. The black 
areas are anthracite and bituminous ; the shaded areas are lignite. (From 
National Geographic Magazine.) 



States now leads the world in coal production, furnishing about one 
third of the total supply. England for many years headed the list, 
and even now furnishes a large amount, for its deposits are extensive. 

Charcoal is a variety of amorphous carbon obtained by 
heating wood, bones, ivory, and other organic matter in 
closed vessels, or by partially burning them in the air. 
There are several varieties. Wood charcoal is a black, 
brittle solid, and often has the shape of the wood from 
which it is made. It is insoluble, though its mineral im- 
purities can be removed by acids. It burns without a flame 
and leaves a white ash (mineral matter). The compact 



CARBON AND ITS OXIDES — CARBIDES 271 

varieties conduct heat and electricity, but porous charcoal 
is a poor conductor. It resists the action of moisture and 
many chemicals ; hence fence posts, telegraph poles, and 
wooden piles are often charred before being put into the 
ground. Most varieties are very porous, and when thrown 
upon water a lump of charcoal floats for some time, owing 
to the presence of air in its pores. Its porosity makes char- 
coal an excellent absorber of gases, some varieties absorb- 
ing ninety times their bulk of ammonia gas. It also absorbs 
colored organic substances from solutions ; this is especially 
true of animal charcoal (see below) . Impure drinking water 
may be partially purified by charcoal, which forms the 
essential part of many filters in houses. Charcoal used for 
such a purpose, however, must be frequently renewed or often 
heated to redness ; otherwise it becomes contaminated. The 
taking up of gases and solids (also liquids) by charcoal is 
called adsorption and is ascribed to the adhesion of the sub- 
stances upon the very large condensing surface of the porous 
charcoal. 

Besides the uses of charcoal mentioned above, it is used 
as a fuel, in the manufacture of steel and of gunpowder, 
and as a medicine. It reduces oxides when heated with 
them, thus : — 



2CuO 


+ c = 


2Cu 


•+ co 2 


Copper 
Oxide 


Carbon 


Copper 


Carbon 
Dioxide 



Wood charcoal is made either in a charcoal pit or kiln, or in a large 
retort. Where wood is plentiful, it is loosely piled into the shape 
shown in Figure 35, and covered with turf to prevent too free access 
of air, though small holes are left at the bottom and a larger one at 
the top (as a central flue), so that sufficient air can pass through the 
pile. The wood is lighted, and as it slowly burns care is taken to 
regulate the supply of air, so that the wood will smolder but not be 
consumed. The volatile matter escapes and the charcoal remains, 
the average yield being about 20 per cent of the weight of the wood. 



272 INORGANIC CHEMISTRY 

This method is crude, uncertain, and wasteful. Much charcoal ia 
now made by heating wood in closed retorts, no air whatever being 
admitted. By this method, which is called dry or destructive dis- 
tillation, the yield of charcoal is 30 per cent, and all the volatile matter 




Fig. 35. — Wood arranged for burning into charcoal. 

is saved. In the ordinary combustion of wood, the hydrogen forms 
water and the oxygen forms carbon dioxide ; but in dry distillation, 
where no oxygen is present, much of the hydrogen forms volatile 
compounds with the carbon and oxygen. Among these volatile 
products are methyl alcohol (CH 4 0), acetone (C 3 H 6 0), and acetic 
acid (C2H4O2). These are commercial substances, and contribute to 
the profit of the process. More or less charcoal is obtained by heating 
any compound of carbon, e.g. sugar or starch, the charring being a 
test for carbon. Such charcoal, especially that obtained from sugar, 
is very pure. 

Animal Charcoal or Bone Black is made by heating bones 
in a closed vessel, or by heating a mixture of blood and 
sodium carbonate. It contains only about 10 per cent of 
carbon distributed throughout the porous mineral matter 
(calcium phosphate) of the bone. Under the name of ivory 
black, animal charcoal is used as a pigment, especially in 
making shoe-blacking. Bone black is extensively used to 
decolorize sugar sirups, oils, and other liquids colored by 
organic matter. (See Adsorption, above.) 



CARBON AND ITS OXIDES — CARBIDES 273 

Coke is made by expelling the volatile matter from bitumi- 
nous coal, somewhat as charcoal is made from wood. It is 
left in the retorts when coal is distilled in the manufacture 
of illuminating gas. On a large scale it is made by heating 
a special grade of bituminous coal in huge ovens, often shaped 
like a beehive, from which air is excluded soon after com- 
bustion begins. Sometimes the coke is made in closed 
retorts constructed to save the by-products, — tar, am- 
monia (in the form of ammonium sulphate or ammoniacal 
liquor), organic compounds (such as benzene, phenol, and 
pyridine), and combustible gases. This method not only 
yields more coke, but is also more profitable because the 
by-products are sold and the combustible gas is used as a 
source of heat, light, and power. Coke is a grayish, porous, 
coherent solid, harder and heavier than charcoal. It burns 
with no smoke and a feeble flame. It contains about 90 
per cent of carbon, the rest being the mineral matter origi- 
nally in the coal. Immense quantities of coke are used in 
the manufacture of iron and steel. It is superior to coal 
for this purpose, because it gives a greater heat when burned, 
reduces oxides easily, and contains little or no sulphur or 
other substances which would impair the metallic product. 
The calorific value of Connellsville coke is about 7900 calories 
per gram. Coke is the fuel used in manufacturing much of 
the pig iron in the United States, and over twelve million 
tons (or about three fourths of the total amount) are made 
annually in the Connellsville district, near Pittsburg, Penn- 
sylvania. 

Gas Carbon is the variety of amorphous carbon which is 
gradually deposited upon the inside of the retorts during 
the manufacture of illuminating gas. It is a black, heavy, 
hard solid, and is almost pure carbon. It is a good con- 
ductor of electricity, and is utilized to some extent for the 



274 INORGANIC CHEMISTRY 

manufacture of the carbon electrodes used in electric lights 
and electric batteries. 

Lampblack is prepared by burning certain oils and resin- 
ous substances rich in carbon in a limited supply of air. 
The dense smoke, which is mainly finely divided carbon, is 
passed through a series of condensing chambers, where it 
is collected upon coarse cloth or a cold surface. Its forma- 
tion is illustrated on a small scale by a smoking lamp, the 
deposited soot being substantially the same as lampblack. 
Lampblack is one of the purest forms of amorphous carbon, 
and is used in making printer's ink and certain black paints. 

Allotropism. — Diamond, graphite, and pure amorphous 
carbon {e.g. sugar charcoal), though exhibiting essentially 
different properties, are identical in composition. All are 
carbon. They can be changed into one another without 
leaving a residue, the amorphous form into graphite and 
finally into diamond, and the diamond itself into amorphous 
carbon. Each burns in oxygen, and the sole product is 
carbon dioxide. Furthermore, a given weight of each 
yields the same weight of carbon dioxide, e.g. when 12 gm. 
of each are burned, 44 gm. of carbon dioxide are always 
produced. There is no doubt about their identity. Ele- 
ments which exist in two or more distinct varieties are called 
allotropic, and the property of assuming more than one 
variety is called allotropism or allotropy (from Greek words 
meaning "another form"). One variety is called an allo- 
trope or an allotropic modification of the other. Allotropism 
is believed by many to be due to a difference in the number 
of atoms in a molecule of the element, and by others to a 
difference in the amount of chemical energy in the allotrope. 

Chemical Properties of Carbon. — Carbon unites directly 
with many elements, though the temperature must usually 



CARBON AND ITS OXIDES — CARBIDES 275 

be raised to bring about combination. Fluorine is the 
only element with which it unites at ordinary temperatures. 
The formation of acetylene (C 2 H 2 ) from carbon and hydrogen 
takes place at the temperature of the electric arc. Carbon 
unites with many metals and non-metals at the high tem- 
peratures produced in the electric furnace, thereby forming 
carbides. (See end of this chapter.) It also unites directly 
with sulphur to form carbon disulphide (CS 2 ) in a special 
form of furnace. A conspicuous and important chemical 
property of carbon is its behavior with oxygen. It unites 
directly with oxygen at elevated temperatures to form car- 
bon dioxide (C0 2 ) or carbon monoxide (CO), depending 
upon conditions w T hich are discussed under these gases. 
Considerable heat is developed during the uniting with 
oxygen. The use of various forms of carbon and carbona- 
ceous substances as fuel is based on this fact. The thermal 
equation for the transformation of carbon (in the form of 
charcoal) into carbon dioxide is — 

C + 2 = C0 2 + 97,000 cal. 

Carbon Oxygen Carbon 

Dioxide 

Carbon reduces oxides, a simple illustration being the re- 
duction of copper oxide, thus : — 



2CuO 


+ c = 


co 2 


+ 


2Cu 


Copper 
Oxide 


Carbon 


Carbon 
Dioxide 




Copper 



This property is applied in the reduction of ores of iron and 
other metals. 

Oxides of Carbon 

Carbon and oxygen do not unite at the ordinary tem- 
perature. But when carbon is heated in air, in oxygen, 
or with some oxides, carbon dioxide (C0 2 ) is usually formed; 



276 INORGANIC CHEMISTRY 

if the supply of oxygen is limited, carbon monoxide (CO) 
is also formed. 

Occurrence and Formation of Carbon Dioxide. — The 

occurrence of carbon dioxide in the atmosphere (3 to 4 
parts in 10,000) and in many natural waters (especially 
mineral waters) has already been mentioned. It is one 
product of ordinary combustion, respiration of animals, 
fermentation, and decay; in all these processes the carbon 
comes from organic matter, while the oxygen comes for the 
most part from the air, though some is supplied by the 
organic matter. 

Ordinary combustion is the chemical combining of carbon 
and oxygen. Hence, when carbon or a substance containing 
it is burned in an excess of air or oxygen, carbon dioxide is 
formed. The equation for this chemical change is — 

C + 2 = C0 2 

Carbon Oxygen Carbon 

Dioxide 

Carbon dioxide is formed, therefore, by the combustion of 
such common carbonaceous substances as wood, coal, 
charcoal, coke, oils, waxes, cotton, bone, starch, sugar, meat, 
bread, alcohol, camphor, and illuminating gas. 

The continuous oxidation of the tissues of the body pro- 
duces carbon dioxide. (See Relation of Oxygen to Life.) 
And if we exhale the breath through a glass tube into calcium 
hydroxide (lime water), the carbon dioxide which is in the 
breath turns the limewater cloudy or turbid — the usual 
test for carbon dioxide. The equation for the change is — 

C0 2 + Ca(OH) 2 = CaC0 3 + H 2 

Carbon Calcium Calcium Water 

Dioxide Hydroxide Carbonate 

When vegetable and animal matter decays, carbon dioxide 
is one product. Many kinds of organic matter ferment, 



CARBON AND ITS OXIDES — CARBIDES 



277 



especially those containing certain kinds of sugar. By 
alcoholic fermentation the sugar changes into carbon dioxide 
and alcohol (see Alcohol), thus: — 

C fl H u O« = 2C0 2 + 2C 2 H 6 

Sugar Carbon Alcohol 



The Preparation of Carbon Dioxide is accomplished in 
the laboratory by the interaction of a carbonate and an acid. 
Calcium carbonate (limestone or marble) and dilute hydro- 
chloric acid are usually used. The equation for the chemical 
change is — 

+ H 2 

Water 

This gas may also be prepared by burning matter contain- 
ing carbon and by strongly heating carbonates (as in making 
lime), thus: — 



CaCO s 


+ 2 HCl = 


co 2 


+ 


CaCl 2 


Calcium 
Carbonate 


Hydrochloric 
Acid 


Carbon 
Dioxide 




Calcium 
Chloride 



C 6 H 12 6 + 6 2 = 


= 6C0 2 


+ 


6H 2 


Sugar Oxygen 


Carbon 
Dioxide 




Water 


CaC0 3 


co 2 


+ 


CaO 


Calcium 
Carbonate 


Carbon 
Dioxide 




Lime 



Properties of Carbon Dioxide. — This gas has many im- 
portant properties besides those mentioned under The 
Atmosphere. It has a slight taste and odor, but no color. 
It is one and a half times heavier than air, and a liter under 
standard conditions weighs 1.977 gm. On account of its 
weight it is usually collected by the displacement of air 
(like chlorine and hydrochloric acid gas). Its weight is 
also one reason why it collects at the bottom of abandoned 
or deep wells, in some valleys near volcanoes, and in mines 
after explosions. At the ordinary temperature and pressure, 
water dissolves its own volume of carbon dioxide. Under 



278 INORGANIC CHEMISTRY 

increased pressure more gas dissolves, which escapes to some 
extent when the pressure is removed. (See Carbonic Acid, 
below.) Hence " soda water/' which is made by forcing 
carbon dioxide into water, effervesces and froths when drawn 
from the soda fountain. Many natural waters and manu- 
factured beverages (such as champagne) sparkle and ef- 
fervesce for the same reason. The critical temperature 
is 31° C. and the critical pressure is about 72 atmospheres. 
Hence this gas can be readily liquefied. It was first liquefied 
by Faraday by the method similar to that used for chlorine. 
Liquid carbon dioxide is now made in large quantities by 
forcing the gas into steel cylinders (by powerful pumps) which 
are cooled by water during the operation; the gas is some- 
times obtained from the fermenting vats of breweries, though 
often prepared by decomposing magnesium or calcium car- 
bonate. When liquid carbon dioxide is allowed to escape 
into the air, a portion evaporates quickly and thereby with- 
draws heat from the remainder; if sufficiently cooled, it 
becomes white, snowlike, solid carbon dioxide. The latter 
is used to produce low temperatures ; a paste of ether and 
solid carbon dioxide has a temperature of — 80° C, and in 
a vacuum it may fall to — 100° C. 

Carbon dioxide does not burn, but extinguishes many 
burning substances, such as a blazing stick or lighted candle, 
the latter being used occasionally to detect carbon dioxide 
in wells, caves, and mines. Air containing from 2.5 to 4 
per cent of carbon dioxide will extinguish small flames. 
Hence the gas is sometimes used to extinguish fires. A stream 
of gas forced upon a small blaze will often prevent a serious 
fire. In the portable fire extinguishers and chemical engines, 
the carbon dioxide is generated rapidly by the interaction 
of sulphuric acid and dissolved sodium bicarbonate (HNaC0 3 ), 
and the pressure of the generated gas forces the saturated 
solution of carbon dioxide out of the extinguisher. 



CARBON AND ITS OXIDES — CARBIDES 279 

Carbon dioxide is reduced by heated carbon into carbon 
monoxide, thus : — 

C0 2 + C = 2 CO 

Carbon Carbon Carbon 

Dioxide Monoxide 

Several metals also decompose carbon dioxide. With 
moderately heated zinc the chemical change is expressed 
thus : — 

C0 2 + Zn = ZnO + CO 

Carbon Zinc Zinc Carbon 

Dioxide Oxide Monoxide 

With burning magnesium, sodium, or potassium the chemical 
change is as follows : — 

C0 2 + 2Mg = 2MgO + C 

Carbon Magnesium Magnesium Carbon 

Dioxide Oxide 

Carbon dioxide combines directly with several oxides, thus : — 
C0 2 + CaO = CaC0 3 

Carbon Calcium Calcium 

Dioxide Oxide Carbonate 

It also unites with water to form carbonic acid (H 2 C0 3 ) 
(see below). 

Relation of Carbon Dioxide to Life. — Animals die when 
put into carbon dioxide. It cuts off the supply of oxygen 
just as water does from a drowning man. The presence 
of a small quantity in the air is objectionable. On the other 
hand, carbon dioxide is an essential food of plants. Through 
their leaves and other green parts they absorb carbon 
dioxide from the atmosphere, decompose it, reject part of 
the oxygen, and store up the carbon as starch ((C 6 Hio0 5 )n). 
The sunlight and the green matter (chlorophyll) aid the plant 
in making its food out of the water (obtained through 
the roots from the soil) and the carbon of the carbon dioxide 
obtained from air. Plants in this way help keep the atmos- 
phere free from an excess of carbon dioxide. 



280 INORGANIC CHEMISTRY 

Carbonic Acid. — Carbon dioxide gas is often called 
carbonic acid gas, or simply carbonic acid. The latter term 
when applied to carbon dioxide (C0 2 ) is incorrect. When 
carbon dioxide is passed into water, it combines to a slight 
extent with the water and forms a weak, unstable acid, 
which is, strictly speaking, carbonic acid. The equation 
for this change is — 



co 2 


+ 


H 2 = 


= H 2 C0 3 


Carbon 




Water 


Carbonic 


Dioxide 






Acid 



Such a solution reddens blue litmus and decolorizes pink 
phenolphthalein, though its action on these indicators is 
rather feeble. In terms of the theory of electrolytic dis- 
sociation it is a weak acid, i.e. it ionizes only to a small 
degree, the ions being H + and HC0 3 ~ (.17 per cent in tenth 
normal solution at 18° C). Carbonic acid is unstable and 
easily breaks up by gentle heat into carbon dioxide and 
water, thus : — 

H 2 C0 3 = C0 2 + H 2 

Carbon dioxide is sometimes called carbonic anhydride, to 
denote its relation to the acid. Carbonic acid is dibasic 
and forms two classes of salts — normal and acid carbonates. 

Carbonates are salts corresponding to carbonic acid. They 
are very common substances. The most abundant natural 
carbonates are those of calcium, magnesium, and iron. 
Immense quantities of sodium and potassium carbonates 
are manufactured. 

A few carbonates are formed by direct combination of an 
oxide and carbon dioxide, but many are formed by the in- 
teraction of carbon dioxide and a base, thus : — 

C0 2 + Ca(OH) 2 = CaC0 3 + H 2 

Calcium Calcium 

Hydroxide Carbonate 



CARBON AND ITS OXIDES — CARBIDES 281 

Others are formed by the interaction of a soluble carbonate 
and another soluble salt, and are readily produced, since all 
except those of sodium, potassium, and ammonium, and 
certain acid carbonates are practically insoluble in water. 
There are two classes of carbonates, the normal and the 
acid. Normal sodium carbonate is Na 2 C0 3 , and acid sodium 
carbonate is HNaC0 3 . The latter is called sodium bicar- 
bonate. Normal calcium carbonate is CaC0 3 , and cal- 
cium bicarbonate or acid calcium carbonate is H 2 Ca(C0 3 ) 2 ; 
the latter salt is unstable, and is decomposed by gentle 
heat into normal calcium carbonate. The acid calcium 
carbonate is formed from the normal carbonate by an excess 
of carbon dioxide. When carbon dioxide is passed into 
water containing the insoluble normal calcium carbonate in 
suspension, the soluble acid calcium carbonate is formed, 
thus: — 

H 2 C0 3 + CaC0 3 = H 2 Ca(C0 3 ) 2 

Carbonic Calcium Acid Calcium 

Acid Carbonate Carbonate 

Now when this solution of acid calcium carbonate is heated, 
the decomposition takes place thus : — 

H 2 Ca(C0 3 ) 2 = CaC0 3 + C0 2 + H 2 

Acid Calcium Calcium Carbon Water 

Carbonate Carbonate Dioxide 

Since many underground waters contain carbon dioxide, 
these waters dissolve the limestone (CaC0 3 ) over which they 
pass, forming " hard " water. When the dissolved acid 
calcium carbonate is decomposed by heat or in some other 
way, the calcium carbonate is reprecipitated. (See Stalac- 
tites (under Calcium Carbonate), Natural Waters, and 
Hardness of Water.) 

Composition of Carbon Dioxide. — If a known weight of pure 
carbon, such as diamond or graphite, is burned in oxygen, it is found 
that for 12 parts of carbon used there are 44 parts of carbon dioxide 
formed. Hence 12 parts of carbon unite with 32 parts of oxygen. 



282 



INORGANIC CHEMISTRY 



Since the vapor density of the gas is 1.375 (if O2 = 32), the molecular 
weight must be 44. These facts necessitate the formula CO2. 

History of Carbon Dioxide. — This gas was described in the seven- 
teenth century by Van Helmont, who called it gas sylvestre. He pre- 
pared it by the interaction of acids and carbonates, detected it in 
mineral water, and observed its formation during combustion and 
fermentation, as well as its action on animals and flames. Black, in 
1755, showed that carbon dioxide is essentially different from ordinary 
air, and that the gas is readily obtained from magnesium and calcium 
carbonates. Since the gas was combined or "fixed" in these sub- 
stances, he called the gas fixed air. His work was verified in 1774 by 
Bergman, who called the gas acid of air. Lavoisier first proved it 
to be an oxide of carbon. 

Carbon Monoxide is formed when carbon is burned in a 
limited supply of air, thus: — ■ 

2C + 2 = 2 CO 

Carbon Oxygen Carbon Monoxide 

If carbon dioxide is passed over heated charcoal, the prod- 
uct is carbon monoxide. That is, carbon reduces carbon 
dioxide to carbon monoxide, the equation for the change 
being — 

C0 2 + C = 2 CO 

Carbon Dioxide Carbon Carbon Monoxide 

These chemical changes take place in every coal fire (Fig. 
36). The oxygen (of the air) entering at the bottom of the 



J t L 



t,a,uh 



A^Wi^\ 




2 CO + o 2 = 2 co 2 



C0 2 + C = 2 CO 



2 C + 2 2 = 2 C0 2 



Fig. 36. — Essential chemical changes during combustion in a coal fire. 



CARBON AND ITS OXIDES — CARBIDES 283 

fire unites with the carbon to form carbon dioxide ; the lattei 
gas in passing through the hot carbon of the fire is reduced 
to carbon monoxide. Some of the carbon monoxide escapes, 
while some burns with a flickering bluish flame on the top 
of the fire, and forms carbon dioxide. 

If steam is passed through a hard coal or coke fire, carbon 
monoxide and hydrogen are formed ; this mixture enriched 
by oil vapors is called water gas and is used as an illuminant 
(page 299). If steam and air are passed through a carbon 
fire, the gaseous product also contains nitrogen and carbon 
dioxide ; this mixture is called producer gas and is used as a 
fuel in industrial processes. 

Preparation of Carbon Monoxide. — This gas is usually 
prepared in the laboratory by gently heating a mixture of 
oxalic acid and sulphuric acid, and collecting the gaseous 
product over water. The oxalic acid decomposes thus : — 

C 2 H 2 4 CO + C0 2 + H 2 

Oxalic Carbon Carbon "Water 

Acid Monoxide Dioxide 

The carbon dioxide can be removed by passing the mixed 
gases through a solution of sodium hydroxide. 

Properties of Carbon Monoxide. — It is a gas without 
color, odor, or taste, and is only slightly soluble in water. 
It burns with a bluish flame, forming carbon dioxide, thus : — 

2 CO + 2 = 2C0 2 

Carbon ' Carbon 

Monoxide Dioxide 

Carbon monoxide is extremely poisonous, and it is very 
dangerous because the lack of odor prevents its detection 
in time to escape its stupefying effect. Many deaths have 
been caused by breathing air containing it. Carbon mo- 
noxide forms a rather stable compound with the haemoglobin 
of the blood, and persons who have been poisoned by it 



284 INORGANIC CHEMISTRY 

cannot usually be revived by air, as in the case of suffoca- 
tion by carbon dioxide. It is an ingredient of water gas 
and most ordinary illuminating gas. Care should always 
be taken to prevent the escape of illuminating gas (as well 
as the gas from a coal stove or furnace) into occupied rooms. 
At a high temperature carbon monoxide readily removes 
combined oxygen, and is, therefore, an important reducing 
agent, e.g. in the manufacture of iron from iron ores in the 
blast furnace. This chemical change may be represented 
thus : — 



Fe 2 3 


+ 


3 CO = 


2Fe 


+ 


3C0 2 


Iron 
Oxide 




Carbon 
Monoxide 


Iron 




Carbon 
Dioxide 



Carbon monoxide, which is sometimes called carbonic 
oxide, forms no acid, and therefore no salts. It does not 
turn calcium hydroxide (limewater) milky, thus being 
readily distinguished from carbon dioxide. Its blue flame 
distinguishes it from all other gases which burn. It unites 
directly with chlorine, to form carbonyl chloride (phosgene, 
COCl 2 ), and with some metals, forming metallic carbonyls, 
e.g. nickel carbonyl (Ni(CO) 4 ). 

Carbides 

Carbides are compounds of carbon and certain elements, 
especially metals. They are numerous, owing to their ready 
formation when carbon and oxides are heated in an electric 
furnace. Calcium carbide and silicon carbide are impor- 
tant commercial substances. Several carbides yield com- 
pounds of carbon and hydrogen by interaction with water, 
acetylene (C 2 H 2 ) and methane (CH 4 ) being commonly 
formed. 

Calcium carbide, CaC 2 , is made on a large scale by heating 
a mixture of lime and coke (a form of carbon) in an electric 



CARBON AND ITS OXIDES — CARBIDES 



285 



furnace. The chemical 


change may be represented 


thus : — 




3 C + CaO 


= CaC 2 + CO 


Carbon Lime 


Calcium Carbon 




Carbide Monoxide 



The furnace in which calcium carbide is made is sketched 
in Fig. 36 a. The mixture of coke and lime (shown in the 
furnace) is intro- 



^VIL 




duced through the 
trap cover A and 
slowly sinks down 
into the space where 
the intense heat is 
produced by the elec- 
tricity as it passes 
between the elec- 
trodes G and E, E. 
The liquid calcium 
carbide is drawn off 
through F. The car- 
bon monoxide rises 
through the pipes D, 
D and enters the up- 
per part of the furnace together with air supplied through 
C, C; this mixture burns and heats the coke and lime. 
The waste gases escape through B. 

Calcium carbide is a hard, brittle, dark gray, crystalline 
solid. Its specific gravity is about 2.2. With water calcium 
carbide forms acetylene (page 290). Owing to its vigorous 
and rapid reaction with water, calcium carbide is packed and 
sold in air-tight cans. 

Silicon carbide, SiC, is a very hard, dark-colored crystalline 
solid prepared by heating coke and sand (silicon dioxide) 
in an electric furnace. (See Carborundum.) 



Fig. 36 a. 



-Electric furnace for making cal- 
cium carbide. 



286 INORGANIC CHEMISTRY 



Problems and Exercises 

1. The specific gravity of charcoal is about 1.5. Why does it 
float on water? 

2. How many grams of calcium carbonate are needed to pre- 
pare 132 gm. of carbon dioxide? 

3. Ten tons of coke were burned, and only 35 tons of carbon 
dioxide were produced. Calculate the per cent of carbon in the 
coke. 

4. Ten grams of carbon dioxide were prepared by the inter- 
action of marble and hydrochloric acid. How many cubic centi- 
meters of marble (specific gravity 2.65) were used? 

5. Suppose 17 gm. of carbon are completely and freely burned 
in air containing 21 per cent of oxygen by volume. Calculate 
(a) the volume of air needed, and (b) the volume and the weight 
of the product. 

6. A mass of limestone is completely transformed by heat into 
a solid and a gas — the usual products. The gas measures 2000 1. 
(standard conditions). Calculate the weight of the three sub- 
stances involved in the reaction. 

7. Calculate the simplest formulas from the following data: 
(a) C = 74.07, N = 17.29, H = 8.64 ; (6) C = 20, O = 26.6, S = 53.3 ; 

(c) 3 1. of an oxide of carbon weigh 3.78 gm. (approximately) ; 

(d) C = 27.27, O = 72.72 ; (e) an oxide of carbon contains 42.857 
per cent of carbon. 

8. A cylindrical tank holds 250 gm. of oxygen. What weight 
and what volume of (a) carbon dioxide and (b) carbon monoxide 
will it hold? (Standard conditions.) 

9. Carbon dioxide is heated with 40 gm. of carbon. What is 
(i) the weight and (b) the volume of the product? (Standard 
conditions.) 

10. Suppose 75 gm. of carbon dioxide are passed slowly over hot 
carbon. Calculate (a) the weight of carbon used and (b) the vol- 
ume of the gaseous product. (Standard conditions.) 

11. Ten grams of oxalic acid are decomposed. What weight of 
carbon monoxide is formed? What volume of carbon monoxide 
at 21° C. and 762 mm.? 

12. What weight of carbon (97 per cent pure) is needed to reduce 
60 gm. of carbon dioxide to carbon monoxide ? What volume of air 
(containing 21 per cent of oxygen by volume) at 19° C. and 758 mm. 
is needed to change the carbon monoxide to carbon dioxide? 



CARBON AND ITS OXIDES — CARBIDES 287 

13. If one volume of carbon monoxide and two volumes of oxy- 
gen are mixed and exploded in a closed space, what will be the vol- 
ume of the resulting gas or gases at the original temperature and 
pressure ? 

14. A ton of calcium carbide is needed. What weight of lime 
and coke must be used ? 

15. Write the formulas of the normal and the acid carbonates 
of Ca, copper, Pb, potassium, Ra, silver, Sr, zinc. (Use Valence 
Tables.) 

16. Starting with carbon, how would you prepare successively 

(a) carbon dioxide, (6) calcium carbonate, (c) lime, (d) calcium 
hydroxide, (e) calcium chloride? 

17. Suggest a method of (a) obtaining carbon from carbon 
dioxide, (6) showing that coal contains mineral matter, and (c) test- 
ing a cave or abandoned mine for carbon dioxide. 

18. State and explain the various chemical changes which occur 
from the entrance of oxygen (in the air) below the grate of a red- 
hot coal fire to the end of the burning of the carbon monoxide at 
the top of the coal. 

19. Express the following reactions by equations : (a) Potassium 
hydroxide and carbon dioxide form potassium carbonate and water, 

(b) barium hydroxide and carbon dioxide form barium carbonate and 
water. 

20. Complete and balance the following equations : (a) BaC0 3 

+ = BaCl 2 + + H 2 0; (b) HNaCOs = Na^CO, + '+ 

C0 2 . 

21. Express the following reactions by volumetric equations: 
(a) Carbon monoxide and oxygen form carbon dioxide ; (6) carbon 
and water (vapor) form hydrogen and carbon monoxide ; (c) car- 
bon dioxide and carbon form carbon monoxide ; (d) carbon and 
oxygen form carbon monoxide ; (e) carbon and oxygen form carbon 
dioxide. 

22. Write the formulas of the following compounds by apply- 
ing the principle of valence: Cadmium carbonate, zinc carbonate, 
barium carbonate, ferrous carbonate. Calculate the weight and 
the volume (standard conditions) of the gas liberated by the inter- 
action of an acid and 25 gm. of each carbonate. 



CHAPTER XVI 

Hydrocarbons — Methane — Ethylene — Acetylene — Illu- 
minating Gas — Flame — Bunsen Burner — Oxidizing 
and Reducing Flames 

Hydrocarbons are compounds of carbon and hydrogen. 
They number about two hundred and fifty.' Their properties 
as individuals vary between wide limits, but as a class the 
hydrocarbons are rather indifferent in chemical behavior. 
They are found in petroleum and its products (kerosene, 
naphtha, lubricating oils, paraffin wax, etc.), in coal tar, in 
coal gas and natural gas, and in some essential oils, such 
as turpentine. On a large scale they are prepared by the 
distillation of petroleum, wood, coal, and coal tar. Indirectly 
the hydrocarbons are the source of many other compounds 
of carbon, which are extensively used in numerous industries; 
several of these organic compounds, as they are called, will 
be described in the next chapter. 

Methane, CH 4 , is the simplest hydrocarbon. It is found 
in coal mines, being a gaseous product of the processes 
which changed vegetable matter into coal. It is called fire 
damp by miners. It is also formed in marshy places by 
the decay of vegetable matter under water, and is therefore 
often called marsh gas. It forms about 90 per cent of 
natural gas, and approximately 35 per cent of the illuminating 
gas obtained by heating coal. 

Methane is formed by catalysis when hydrogen is passed 
over a mixture of carbon and nickel at 250° C. It is usually 
prepared by heating a mixture of sodium acetate, sodium 

288 



METHANE AND ETHYLENE 289 

hydroxide, and lime. It may also be prepared by the inter- 
action of aluminium carbide and water, thus : — 



A1 4 C S + 


12H 2 = 


= 3 CH 4 


+ 


4 A1(0H), 


Aluminium 


Water 


Methane 




Aluminium 


Carbide 








Hydroxide 



Methane has no color, taste, or odor. It burns with a 
pale, bluish flame. A mixture of methane with oxygen or 
air explodes violently when ignited by a spark or flame. 
Terrible disasters occur in coal mines from this cause. The 
products of the explosion are carbon dioxide and water, thus : 



CH 4 


+ 2 2 = 


= co 2 + 


2H 2 


letham 


j Oxygen 


Carbon 
Dioxide 


Water 



The carbon dioxide, called choke damp or black damp by 
the miners, often suffocates those who escape from the ex- 
plosion. A liter of methane weighs ,717 gm. (standard 
conditions) . 

Ethylene, C 2 H 4 , or olefiant gas, is formed by the destructive 
distillation of wood and coal. It is usually prepared in the 
laboratory by heating a mixture of concentrated sulphuric 
acid and ethyl alcohol (i.e. ordinary alcohol), and collecting 
the gas over water. The essential change is represented 
thus : — 

C 2 H 6 = C 2 H 4 + H 2 

Alcohol Ethylene 

Ethylene is a colorless gas and has a pleasant odor. Its 
critical temperature is 9° C. and its critical pressure is 58 
atmospheres, so it can be readily condensed to a liquid. 
The latter by evaporation in a vacuum produces a tempera- 
ture as low as — 140° C. A liter of ethylene weighs 1.25 gm. 
(standard conditions). The gas burns with a yellow flame, 
and is one of the essential illuminating ingredients of coal gas. 
When ethylene burns, the complete combustion is represented 
thus : — 



290 INORGANIC CHEMISTRY 



2H4 + 


3< 


3 2 


= 2C0 2 


+ 


2H 2 


lylene 






Carbon 
Dioxide 




Water 



If mixed with oxygen in this proportion and ignited, the mix- 
ture explodes. Ethylene reduces potassium permanganate. 

Acetylene, C 2 H 2 , is formed by the direct union of hydro- 
gen and carbon when an electric arc is produced between 
two carbon rods in hydrogen gas. This method of forma- 
tion, though not convenient, is interesting, because very few 
hydrocarbons have as yet been formed directly from their 
elements. A small quantity is present in coal gas. It is also 
formed by the incomplete combustion of coal gas, e.g. when 
the flame of a Bunsen burner strikes back and burns at the 
base. (See Bunsen Burner, below.) Acetylene is now pre- 
pared cheaply on a large scale by the interaction of calcium 
carbide and water, thus : — 



CaC 2 + 


2H 2 


= C 2 H 2 + 


Ca(OH) 2 


Calcium 


Water 


Acetylene 


Calcium 


Carbide 






Hydroxide 



Acetylene is a colorless gas, and, if impure, has an offen- 
sive odor. It is poisonous if breathed in large quantities, 
but much less dangerous than gases like carbon monoxide. 
Its density is about .92. A liter weighs 1.162 gm. (stand, 
cond.). Water at the ordinary temperature dissolves its own 
volume of the gas. Reliable tests show that acetylene 
does not act upon any common metal or alloy, though it 
forms explosive compounds with salts of metals, especially 
of copper. As a precaution, copper and brass are seldom 
used in large vessels containing or generating acetylene, 
though they might safely be used in small lamps. 

The critical pressure is about 61 atmospheres and the critical tem- 
perature is about 35° C, so acetylene can be readily liquefied. 
Cylinders of liquid acetylene have exploded, causing loss of life and 
destruction of property, and its use in this form has been prohibited 



ACETYLENE 



291 



in some localities. Under ordinary atmospheric pressure acetylene 
gas does not explode, but if subjected to a pressure of two or more 
atmospheres, it will explode by a shock or when a spark or flame is 
brought near it. A mixture of acetylene and air, if ignited, explodes. 
The mixture to be explosive, however, must contain a large per cent 
of acetylene (a condition hardly possible except from sheer careless- 
ness, because the disagreeable odor reveals the presence of the gas). 
Acetylene gas, it is evident, must be used with the same precaution 
as any ether illuminating gas. 

Acetylene is found by analysis to contain only carbon and hydro- 
gen combined in the ratio of 12 to 1 by weight. Its vapor density is 
.8125 (if 2 = 32). Therefore its molecular weight is 26 and its for- 
mula is C2H2. The graphic formula of acetylene is usually written 
thus: H-C = C-H. 

Acetylene as an Illuminant. — Acetylene burns in the air 
with a luminous, smoky flame. But when air is mixed 
with the gas as the latter 
issues from a small opening, 
the mixture burns with a bril- 
liant, white flame, w 7 hich does 
not smoke. It is extensively 
used as an illuminant. The 
flame is almost like sunlight, 
hence by the acetylene flame 

most Colors appear the Same Fig. 37. —Relative size of acetylene 




and illuminating gas flames giving 
the same amount of light. The 
acetylene (smaller) flame con- 
sumes only one tenth as much gas 
an hour as the illuminating gas 
flame. (One half actual size.) 



as in daylight. It is adapted 
for taking photographs, since 
its action closely resembles 
that of the sun. It is a dif- 
fusive light, and the flame is 
much smaller than an ordinary gas flame of the same 
lighting power (Fig. 37). 

With a proper burner the combustion of acetylene is complete 
and may be represented thus : — 



2 CoH 2 + 5 Oo = 4 C0 2 + 2 H 8 



292 INORGANIC CHEMISTRY 

Ordinary gas burners cannot be used for acetylene. In acety- 
lene burners the gas issues from two small holes drilled at an angle, 
so that the jets strike each other and produce a flat flame (Fig. 38). 





I 

Fig. 38. — Acetylene flame. Fig. 39. — Acetylene burner. 

Other holes, such as B, B, B, B, permit air to be drawn in mechani- 
cally by the acetylene as it rushes through the burner. The open- 
ings for the mixture are so fine that the flame cannot strike back 
and cause an explosion (Fig. 39). 

Generation of Acetylene. — The ease with which acetylene is gen- 
erated can be shown by putting a little water in a test tube and then 
dropping in a small lump of calcium carbide. The gas bubbles 
through the liquid ; after the action has gone on long enough to expel 
the air, the acetylene may be lighted by holding a burning match at 
the mouth of the tube. On a larger scale, the gas can be generated 
by putting the calcium carbide into a flask provided with a dropping 
funnel and delivery tube, and allowing water to drop slowly upon the 
carbide; the gas thus generated can be collected in bottles over 
water. There are two classes of commercial generators. In one, 
water is added to the carbide, but in the other the carbide drops into 
the water. The intense heat liberated when calcium carbide inter- 
acts with water decomposes acetylene ; hence, a generator to be 
effective and safe should be constructed so that this heat will be ab- 
sorbed. The first class of generators is dangerous, except when a 
small quantity of gas is desired, as on the lecture table or in a lantern. 
In the second class, a small amount of calcium carbide drops automat- 
ically into a large volume of water as fast as the gas is needed, thus 
insuring a pure, cool gas, and eliminating the danger of an explosion. 
A pound of calcium carbide yields about five cubic feet of acetylene 
gas. Acetylene dissolves in acetone, and cylinders in which con- 
siderable gas has been dissolved under pressure are used to furnish 
gas for automobile lanterns. 

Acetylene is an endothermic compound, and several of its proper- 
ties described above are due to its endothermal nature. It is formed 



ACETYLENE 293 

from its elements with absorption of heat. Being endothermie, it 
explodes when subjected to a shock (if the gas is under a pressure of 
two or more atmospheres), thus: — 

C 2 H 2 = 2 C + H 2 + 53,100 cal. 

The temperature of the oxy-acetylene flame is nearly 3000° C. ; 
the thermal equation for the chemical change is — 

2 C 2 H 2 +50 2 =4C0 2 +2 H 2 + 310,000 cal. 

The flame when produced in a suitable burner is used in welding 
metals and in dismantling metal structures, e.g. bridges, frames of 
buildings, abandoned battleships, etc. The burner resembles an 
oxyhydrogen burner (Fig. 4). 

Petroleum is the source of numerous hydrocarbons. It is 
an oily liquid obtained from the earth in many parts of the 
world. In the United States the chief localities are Ohio, New 
York, Pennsylvania, West Virginia, Kentucky, Indiana, 
Colorado, Texas, and California. The immense deposits in 
Russia are in the Baku district on the Caspian Sea. Some 
is also found in Canada, India, Japan, and Austria. 

Crude petroleum is an oily liquid, with an unpleasant 
odor. Its color varies from straw to greenish black, and 
most kinds are greenish in reflected light. It usually 
floats upon water. Its composition is complex, but all 
varieties are essentially mixtures of several liquid and solid 
hydrocarbons. American oils contain chiefly members of 
the paraffin series {i.e. methane series). Some varieties 
contain compounds of nitrogen and of sulphur. 

In some localities the oil issues from the earth, but it is usually 
necessary to drill through rocks and insert a pipe into the porous 
rock containing the oil. At first the oil often "shoots" out of the 
well in tremendous volumes, owing to the pressure of the confined 
gas, but after a time a pump is needed to draw it to the surface. 
The oil is then forced by powerful pumps through large pipes to 
central points for storage or for delivery to refineries, which are often 
many miles from the oil well. This network of pipes in the eastern 
United States is over 25,000 miles long. 



294 INORGANIC CHEMISTRY 

Some crude petroleum is used in making water gas (see 
below) and as fuel on locomotives and steamships, but most 
of it is separated into various commercial products.' This 
process is called refining. The petroleum is cleaned by set- 
tling and filtration, and then distilled in huge iron vessels. 
The vapors are condensed as they pass through coiled pipes 
immersed in cold water. Certain products are obtained 
from the residue left in the still. (See below.) 

The different distillates, which are collected in separate tanks, 
are further separated by redistillation and purified by special treat- 
ment. The commercial products thus obtained are petroleum 
ether (b. p. 40-70° C), gasolene (70-90), naphtha (90-120), ben- 
zine (120-150), kerosene (150-300) ; various grades of these prod- 
ucts are distinguished by boiling points or by specific gravity. 
These liquids are mixtures of several different hydrocarbons. They 
are widely used as solvents, fuels, illuminants, and in making gas. 

Gasolene is used as a fuel in internal combustion engines. The 
vapor of gasolene burns readily ; if the vapor is mixed with air 
and the mixture is ignited properly, the combustion is so rapid that it 
is practically an explosion and the suddenly expanded gas exerts 
a pressure which is converted by the machinery into steady and 
continuous motion. Kerosene is the well-known illuminating oil. 
It is carefully freed from inflammable liquids and gases and from 
tarry matter and semisolid hydrocarbons by agitating it successively 
with sulphuric acid, sodium hydroxide, and water. Commercial 
kerosene must have a legal flashing point. Flashing point is " the 
temperature at which the oil gives off sufficient vapor to form a 
momentary flash when a small flame is brought near its surface.' ' 
The legal flashing point varies in different localities from 44° C. 
to 68° C. 

From the residuum left in the still after the last distillation 
lubricating oils, vaseline, and paraffin wax are obtained by further 
treatment. Mineral lubricating oils have largely replaced animal 
and vegetable oils. Vaseline finds extensive use as an ointment. 
Paraffin wax is used to make candles, to waterproof paper, to extract 
oils from plants and flowers ; also as a coating for many substances, 
thereby producing a smooth surface or facilitating slow combustion 
(as in certain matches). The final residue left in the retorts is coke. 



ILLUMINATING GAS 295 

Hydrocarbons are often extracted from it, some is made into electric 
light carbons, and some is used as a fuel. 

Natural Gas is a combustible gas, which issues from the 
earth in many places, Methane forms about 90 per cent 
of the mixture. It is used as a fuel for heating houses, 
generating steam, and manufacturing iron, steel, glass, 
brick, and pottery. In Ohio, Indiana, and other gas- 
producing regions of the United States, wells, like petroleum 
wells, are drilled for the escape of natural gas, which is 
distributed to consumers through pipes similar to those 
used for illuminating gas. 

Illuminating Gas. — Besides acetylene there are other 
kinds of illuminating gas. Coal gas and water gas are the 
most common. 

Coal Gas is made by distilling bituminous coal and purify- 
ing the volatile product. The hydrogen in the coal passes 
off partly as free hydrogen, and partly in combination with 
carbon as hydrocarbons and with nitrogen as ammonia. 
The ammonia, carbon dioxide, and sulphur compounds 
are regarded as impurities, and are removed as completely 
as possible before the gas is delivered to the consumer. The 
by-products are coke, gas carbon, tar, and ammoniacal 
liquor. The essential parts of a coal gas plant are shown 
in Figure 40. 

The coal is distilled in air-tight O-shaped retorts made of fire 
clay. The volatile products escape through a pipe, and bubble 
through water into the hydraulic main. Here some of the tar is 
deposited and the ammonium compounds are dissolved by the water 
that flows constantly through the main. This water is kept at a 
constant level and acts as a "seal" to prevent the gas from passing 
back into the retorts. The ammoniacal liquor and tar flow into 
separate wells. From the hydraulic main the gas, which is hot and 
impure, passes into the condenser. This is a long series of vertical 
iron pipes, so constructed that the gas must pass through the entire 



296 



INORGANIC CHEMISTRY 



T7 




r / un i J u \v 



siuQijy 



ILLUMINATING GAS 297 

length of the pipes, while the tar and ammoniacal liquor flow into 
the proper receptacles. The main object of the condenser is to 
cool the gas slowly and condense and remove the tar. An exhauster, 
in most plants, draws or forces the gas from the hydraulic main 
through the condenser into the scrubber and onward through the 
purifiers into the gas holder. The exhauster also reduces the pressure 
in the retorts and regulates the pressure in the holder. (See below.) 
From the exhauster the gas passes into the scrubber. Its purpose 
is to remove the remaining ammonia, part of the carbon dioxide 
and the hydrogen sulphide gas, and the last traces of tar. Scrubbers 
vary in construction. One form is a double tower filled with wooden 
slats or with trays covered with coke or pebbles over which ammo- 
niacal liquor slowly trickles in the first part and pure water in the 
second. The gas enters at the bottom, meets the descending liquid, 
and is thoroughly washed. Another form widely used consists of a 
cylindrical vessel in which numerous wooden slats revolve in com- 
partments and dip into ammoniacal liquor or water at the bottom. 
The liquid forms a film on the slats and absorbs the ammonia and 
other gases, while the resulting solution mixes with the liquor at 
the bottom and flows into the proper well. Sometimes a separate 
tar extractor is connected with the scrubber. This is a tower filled 
with perforated plates, which catch and remove the tar mechanically 
as the gas passes through into the scrubber. From the scrubber 
the gas passes into the purifiers. Their chief purpose is to remove 
the remaining carbon dioxide and sulphur compounds. They are 
shallow, rectangular iron boxes provided with slat frames loosely 
covered with lime. In some plants iron oxide is used as the purify- 
ing material. 

The purified gas next passes through a large meter, which records 
its volume, into a gas holder. The holder is an enormous, cylindrical, 
iron tank in which the gas is stored. Weights and the pressure 
from the exhauster so balance it that it exerts just enough pressure 
to force the gas through the pipes to the consumer. 

A ton of good coal yields about 10,000 cubic feet of gas, 1400 
pounds of coke, 120 pounds of tar, 20 gallons of ammoniacal liquor, 
and a varying amount of gas carbon. The coke, which remains in 
the retorts after distillation, is sold as fuel. The tar, or coal tar as 
it is often called, collected from the hydraulic main and condenser, 
is a thick, black, foul-smelling liquid. Some is used for preserving 
timber, making tarred paper and concrete, and as a protective paint. 
Most of it is now separated by distillation into its more important 



298 



INORGANIC CHEMISTRY 




ILLUMINATING GAS 299 

components, especially benzene (CeH 6 ). The ammoniacal liquor 
from the hydraulic main, condenser, and scrubber is the source of 
ammonia and its compounds. Gas carbon is the hard deposit which 
collects on the inside of the retort, and is used in the electrical in- 
dustries. (See Gas Carbon.) 

Water Gas is mainly a mixture of hydrogen and carbon 
monoxide. It is made by forcing steam through a mass of 
hot anthracite coal and mixing the gaseous product with 
hot gases obtained from oil. The essential parts of the 
apparatus are shown in Figure 41. 

Air is forced through the coal fire in the generator, and the hot 
gases which are produced pass down the carbureter, up into the 
superheater, and escape through its top into the open air. This 
operation lasts about four minutes, and is called the " blow." It 
heats the fire brick inside the carbureter and superheater intensely 
hot, air often being forced in to raise the temperature. The air 
valves and the top of the superheater are now closed, and the " run " 
begins, which lasts about six minutes. Steam is forced into the 
generator at the bottom. In passing through the mass of incan- 
descent carbon the steam and carbon interact thus : — 

C + H 2 = CO + H 2 

Carbon Steam Carbon Hydrogen 

Monoxide 

This mixture of hydrogen and carbon monoxide burns with a 
feeble though hot flame, and is sometimes used as fuel. Before this 
mixture can be used as an illuminating gas it must be enriched with 
gases which are illuminants. Therefore, the mixed gases pass to 
the top of the carbureter, where they meet a spray of oil. And as 
the gaseous mixture passes down the carbureter and up the super- 
heater, the hydrocarbons of the oil are transformed by the intense 
heat into hydrocarbons which do not liquefy when the gas is cooled. 
The addition of hydrocarbons is called carbureting. From the 
superheater the water gas passes through the purifying apparatus 
into a holder. 

Oil Gas is an illuminating gas made from petroleum. When 
petroleum is heated under proper conditions, complex reac- 
tions occur which consist largely in the decomposition of the 



300 



INORGANIC CHEMISTRY 



heavier hydrocarbons ; this process is called " cracking. " The 
gaseous product by suitable treatment yields an illuminating 
gas containing a larger proportion (45 per cent) of illuminants 
than coal gas. Pintsch gas is an oil gas. 

Characteristics of Illuminating Gases. — Coal gas and 
water gas have a disagreeable odor. They are mixtures. 
The following table shows the average : — 



Composition of Illuminating Gases 




Constituents 


Coal Gas 


Water Gas 


Ethylene (and other illuminants) .... 
Methane 


5.0 
34.5 
49.0 

7.2 
1.1 
3.2 


16.6 
19.8 


Hydrogen 

Carbon monoxide 


32.1 
26.1 


Carbon dioxide 


3.0 


Nitrogen 


2.4 







Both kinds of illuminating gas may contain a little oxy- 
gen and traces of ammonia and hydrogen sulphide gas. 
Nitrogen and the last portions of carbon dioxide are impu- 
rities not easily removed. Methane, hydrogen, and carbon 
monoxide burn with a feeble (non-yellow) flame, and are 
often called diluents; they furnish heat, but no light. 

The luminosity of illuminating gas depends mainly upon 
the presence of the hydrocarbons that contain a relatively 
large proportion of carbon. The most important illuminants 
in coal gas and water gas are ethylene (and analogous 
hydrocarbons in the same series), acetylene, and benzene 
(C 6 H 6 ). The utility of an illuminating gas depends upon its 
luminosity. This property is measured by a photometer 
and is expressed in " candles," or candle power. The 
determination of candle power is made by comparing the 
light produced by burning the gas in a standard burner 



FLAME 301 

at the rate of five cubic feet an hour with the light pro- 
duced by a standard wax candle burning at the rate of 120 
grains (7.77 gm.) an hour or by a standard flame. Thus, 
a gas flame 20 times brighter than the standard has a can- 
dle power of 2®. The candle power of ordinary coal gas is 
about 17, and that of water gas is about 25. Ordinary illumi- 
nating gas has a candle power of about 20. The candle 
power of oil gas is 50 or more. 

Flame. — Chemical action, as we have already seen, is 
often accompanied by light. Sometimes the light is merely 
a glow, as in the case of burning charcoal, or a shower of 
sparks, seen when iron burns in oxygen, or a mere flash, 
displayed by an explosion of hydrogen and oxygen. But 
when the interacting substances are gases or vapors from 
volatilized liquids or solids, the chemical change is accom- 
panied by a more or less quiet and continuous light. The 
term flame is commonly applied to such a light, though 
a flame is really a series of chemical ^changes in which the 
gases interact at such a temperature that light is produced. 
Popularly, a flame is a gas burning in the air, i.e. the gas 
or some of its elementary constituents are combining with 
the oxygen of the air. But since the flame is due primarily 
to the chemical combination of the oxygen of the air and the 
gas, it is immaterial where the actual change takes place. 
Ordinarily, the gas burns in the surrounding air, but the 
flame is produced just as truly, though not so conveniently, 
in a vessel of illuminating gas, provided, of course, air is 
supplied. This change of atmospheres, so to speak, is shown 
by a simple experiment (Fig. 42). The lamp chimney B 
is filled with illuminating gas through the bent tube D, and 
its escape is temporarily prevented by closing the opening 
in the asbestos cover A. The gas is lighted at the lower 
end of the tube C, and when the hole in A is uncovered, 



302 



INORGANIC CHEMISTRY 



the flame rises in C and continues at the end within the 
chimney as long as air is drawn up through C and gas supplied 
through D. The unconsumed illuminating 
gas escapes through the hole in A, and if 
ignited, burns as shown in the figure. Chem- 
ically both flames are alike. The outer flame 
is in an atmosphere of air, while the inner 
flame is in an atmosphere of illuminating gas; 
but both are due to the combination of oxy- 
gen with the elementary constituents of the 
illuminating gas. 

In an illuminating gas flame the gas itself 
is burning in the air. In a lamp flame the 
burning gas comes from the oil which is drawn 
up through the wdck and then volatilized by 
the heat. Similarly, in a candle flame the 




Fig. 42. — Com- 
bustion in illu- 
minating gas burning gas comes from the melted wax. 
The flame produced by most burning hydro- 
carbons is luminous and has a yellowish white color. 



Luminous Flames. — The luminous hydrocarbon flame 
has several distinct parts, and the structure of the flame is 
essentially the same, whether produced by burning illuminat- 
ing gas, kerosene oil, or candle wax. The candle flame may 
be taken as the type. An examination of the enlarged 
vertical section shown in Figure 43 reveals four somewhat 
conical portions. (1) Around the wick there is a black 
cone (A), filled with combustible gases formed from the 
melted wax. They do not burn because no oxygen is pres- 
ent. With a glass tube of fine bore it is possible to draw 
off these gases from a large flame and light them at the 
upper end of the tube. (2) Around the lower part of the 
dark cone is a faint bluish cup-shaped part (5, B). It 
is the lower portion of the exterior cone where complete 



LUMINOUS FLAMES 



303 



combustion of the gases occurs, since plenty of oxygen from 
the air reaches this portion. (3) Above the dark cone is 
the luminous portion ((7). It is the largest and most im- 
portant part of the flame. It is usually spoken of as "the 





Fig. 43. — Typical candle 
flame. 



Fig. 44. — Paper charred by a candle 

flame. 



flame." Combustion is incomplete here, because little or 
no oxygen can pass through the exterior cone. The tem- 
perature is high, however, and the hydrocarbons undergo 
complex changes. Acetylene is probably formed. The 
most characteristic change is the liberation of small particles 
of carbon. This liberated carbon, heated to incandescence 
by the burning gases, makes the flame luminous. The carbon 
glows but does not burn up, because little or no oxygen is 
present. A piece of crayon or glass rod held in this part 
of the flame is at once coated with soot, which consists of 
very fine particles of carbon. (4) The exterior cone (D, D) 
is almost invisible. Here the combustion is complete, 
because the oxygen of the air changes all the carbon into 
carbon dioxide. That this is the hottest region of the flame 
may be easily shown by pressing a piece of stiff white paper 
for an instant down upon the flame almost to the wick. 



304 INORGANIC CHEMISTRY 

The paper will be charred by the hot outer portion of the 
flame, as shown in Figure 44. 

These four portions may be found in all luminous hydro- 
carbon flames, whatever the shape. An ordinary gas flame 
is flattened by forcing the gas through a narrow slit in the 
burner tip, so that the flame will give more light. The blue 
part is easily seen, however, when the gas flame is turned 
low or looked at through a small opening; the dark and yel- 
low parts are always visible — the latter being intentionally 
enlarged. The flat or circular flame of an oil lamp likewise 
presents the same characteristics. 

The gaseous products of the combustion of hydrocarbons 
are water vapor and carbon dioxide. A bottle in which 
a candle is burning has, at first, a deposit of moisture on 
the inside; and if the candle is removed and lime water added, 
the presence of carbon dioxide is shown by the cloudiness 
of the limewater. The oxygen needed by the burning 
hydrocarbons is obtained from the air. If not enough oxy- 
gen is present, the flame smokes, i.e. the carbon is thrown off 
into the air before the particles are heated hot enough to 
glow. All oil lamps are so constructed that air enters the 
burner below the flame. Large oil lamps have a central 
opening through which a large volume of air passes up inside 
the circular flame. Otherwise the lamp would burn with 
a very smoky flame. 

The luminosity of hydrocarbon flames is affected by other 
factors besides the presence of glowing carbon. One of these 
is temperature. Gases cooled before being burned give 
very poor light. A candle flame may be cooled enough to 
extinguish it. Thus, if a coil of copper wire is lowered upon 
a candle flame, the flame smokes, loses its yellow color, and 
finally goes out; but if a coil of hot wire is used, the flame 
burns unchanged. Gases, as well as solids and liquids, have 
a kindling temperature, i.e. a temperature to which they 



THE BUNSEN BURNER AND ITS FLAME 305 

must be heated before they " catch fire." This temperature 
differs with different substances. As we lower the tem- 
perature of gases burning with a luminous flame, their lu- 
minosity decreases, and below their kindling point they 
will not burn. The density of the gases in the flame and of 
the atmosphere itself likewise modifies luminosity. The 
flame of a candle was found by experiment to be smaller 
on the top of Mont Blanc than at the base. 

Not all luminous flames are hydrocarbon flames. Thus, 
magnesium burns with a brilliant flame. Its luminosity 
is due to the incandescence of solid particles of magnesium 
oxide. Similarly, the bright flame of burning phosphorus 
is accounted for by the incandescent particles of solid 
phosphorus pentoxide. Most luminous flames contain solid 
particles, though in a few cases luminosity is caused by the 
combustion of gases under pressure, no solid particles what- 
ever being produced. 

Non-Luminous Flames. — Not all flames are luminous. 
The hydrogen flame is almost invisible in air and oxygen, 
but pale blue in chlorine. The flames of carbon monoxide 
and methane are also a faint blue. The most common 
non-luminous flame is the Bunsen flame. 

The Bunsen Burner and its Flame. — When illuminating 
gas is mixed with air before burning, and the mixture burned 
in a suitable burner, a flame is produced which is non- 
luminous and very hot. The temperature of the hottest 
part is about 1500° C. This flame deposits no carbon, 
since its products are entirely gaseous. Such a flame is 
called the Bunsen flame, for it was first produced in a burner 
devised by the German chemist Bunsen. This burner is 
constantly used in chemical laboratories as a source of heat, 
and modified forms have numerous uses. One form, for 
example, furnishes the heat in the gas range used for cooking. 



306 



INORGANIC CHEMISTRY 



O 



o 



The parts of a typical Bunsen burner are shown in Figure 45. 
The gas enters the base and escapes through a very small 
opening into the long tube, which screws down 
over this opening. At the lower end of the long 
tube there are two holes, through which air is 
drawn by the gas as it rushes out of the small 
opening. The gas and air mix as they rise in the 
tube, and this mixture of air and gas burns at the 
top of the long tube. The size of the air holes at 
the bottom of the long tube may be changed by 
a movable ring, thus varying the volume of the 
entering air. When the holes are open, the typical 
non-luminous, hot Bunsen flame is formed. The 
combustion of the constituents of the hydrocar- 
bons is practically complete. The non-luminous 
flame is free from soot, therefore apparatus heated 
by this flame is not blackened. The Bunsen flame 
can be made luminous by closing the air holes or 
by introducing fine particles into 
the flame, — such as charcoal dust, 
finely divided metals, and sodium 
compounds. 

It was formerly believed that the 
non-luminous character of the Bun- 
sen flame is solely due to the complete combustion of the 
carbon by the oxygen of the entering air. Recent experi- 
ments have shown, however, that the result is partly due 
to the diluting action of the nitrogen. 

The gas burns at the top of the tube, not inside, because 
the proper mixture of gas and air flows out more quickly 
than the flame can travel back through the tube to the small 
exit. If the gas supply is slowly decreased, the flame be- 
comes smaller, disappears with a slight explosion, and burns 
at the exit inside the tube. A sudden draft of air, too large 




Fig. 45. — Parts of a typical 
Bunsen burner. 



THE BUXSEX BURNER AXD ITS FLAME 307 

holes at the lower end of the tube, or too low gas pressure 
also may cause the flame to " strike back/' as this action is 
called. This change is due to the fact that the tube con- 
tains an explosive mixture of air and illuminating gas, 
through which the flame travels downward faster than the 
mixture escapes from the tube. This modified flame has 
a pale color and disagreeable odor, and deposits soot. 

The Bunsen flame has many characteristic properties. 
Its color is bluish, and the different cones have different 
tints. There are really three cones: (1) the blue or greenish 
inner one of unburned gases; (2) the very faint blue middle 
one; (3) and the outer one, which is a pale blue, and rep- 
resents the blue cone in the candle flame. The middle 
and outer cones are not always easily distinguished; so for 
all practical purposes it is convenient to divide the flame 
into two parts, — an inner cone of unburned gases and an 
outer cone in which all the carbon is consumed. Combustible 
gases may be drawn off by a tube from the inner cone and 
ignited. A match laid for an instant across the top of the 
tube is charred only at the two points where it touches 
the outer cone; and a sulphur match suspended by a pin 
across the top of an unlighted burner is not kindled until 
several minutes after the 
gas is first lighted. 

A piece of wire gauze 
pressed down upon the Bun- 
sen flame shows a dark cen- 
tral portion surrounded by 
a luminous ring. The flame 
is beneath the gauze, al- 
though the gas passes freely FlG - 46> — wire gauze and flame - 
through it and escapes. If the gas is extinguished and then 
relighted above the gauze, it will burn above but not beneath 
(Fig. 46). The gauze cools the gas below its kindling tern- 




D 



308 



INORGANIC CHEMISTRY 




perature. The miner's safety lamp invented by Davy depends 
upon this last principle. It is an oil lamp surrounded by a 
cylinder of fine wire gauze (Fig. 47). When 
taken into a mine where there are explosive 
gases (such as methane — see Methane), the 
flame continues to burn inside, though its 
size and color change. The gas often enters 
the lamp and burns inside, but the flame 
within does not ignite the gases without 
because the wire gauze keeps them cooled 
below their kindling temperature. Hence an 
explosion is often prevented. When miners 
"notice changes in the lamp flame, they usually 

Fig. 47. — One see k a sa f e place, 
form of Davy's 
safety lamp. 

Oxidizing and Reducing Flames. — The outer 
portion of the Bunsen flame is called the oxidizing flame, 
because here the oxygen is freely given to substances. The 
inner portion is called the reducing flame, 
because here the hydrocarbons withdraw 
oxygen. A sketch of the general relation 
of these flames is shown in Figure 48. 
A is the most effective part of the oxidiz- 
ing flame, and B of the reducing flame. 
At A metals are oxidized, and at B oxy- 
gen compounds are reduced. Sometimes 
a long tube with a small opening at one 
end, called a blowpipe, is used to produce 
these flames. Another tube with a flat- 
tened top is put inside the burner to pro- 
duce a luminous flame. The tip of the 
blowpipe rests in or near this flame, and 
if air is gently and continuously blown ^ zing ( 7)\nd°rJ 
through the blowpipe, a long, slender flame ducing (B) flames. 





OXIDIZING AND REDUCING FLAMES 309 

is produced, called a blowpipe flame (Fig. 49). It is like 
the Bunsen flame as far as its oxidizing and reducing prop- 
erties are concerned. The blowpipe is 
used in the laboratory and by jewelers 
and mineralogists. On a large scale the 
blowpipe flame is used to reduce or oxidize 
ores. (See Compound Blowpipe.) 

The Bunsen flame is extensively utilized flame, showing oxi- 
in producing the Welsbach light. The J™^ gJJ ^* re " 
non-luminous flame heats an inverted 
bag or mantle of oxides of thorium and cerium, and the 
mantle glows with an intense light. The candle power 
varies from 40 to 100. The proportion of thorium oxide to 
cerium oxide in the mantle is 99 to 1. 

Problems 

1. Calculate the weight of carbon in (a) 32 gm. of methane, 
(b) 75 gm. of ethylene, and (c) 145 gm. of acetylene. 

2. What volume and what weight of oxygen are needed for the 
complete combustion of (a) 15 gm. of methane, (b) 20 gm. of ethy- 
lene, and (c) 25 gm. of acetylene? (Standard conditions.) 

3. How many grams of potassium chlorate are required to fur- 
nish the oxygen necessary to burn 10 liters of methane, and how 
many liters of each of the products will be formed? (Standard 
conditions.) 

4. A gas holder has a maximum capacity of 12,000 cubic meters. 
How much calcium carbide (92 per cent pure) must be used to fill 
the holder with acetylene gas measured at 20° C. and 757 mm. ? 

5. What weight of acetylene can be prepared from (a) a metric 
ton of pure calcium carbide and (6) a pound of calcium carbide 
which is 90 per cent pure ? 



CHAPTER XVII 
Other Carbon Compounds 

Introduction. — It was formerly believed that starch, sugar, 
and other compounds obtained from plants and animals 
were produced by the influence of some mysterious vital force. 
Such compounds were called organic, because of their con- 
nection with living things, i.e. with bodies having organs; 
and they were sharply distinguished from inorganic or 
mineral compounds obtained from the earth's crust. This 
distinction prevailed until Wohler, in 1828, prepared urea — 
a distinct organic compound — from inorganic substances. 
Since then the barrier between the two classes of compounds 
has been completely removed. It is now believed that com- 
pounds of carbon, whatever their source, are subject to the 
laws that govern all other compounds. The terms organic 
and inorganic are still used, though they have lost their 
original narrow meaning. Carbon forms a vast number of 
compounds which are related to each other, and which differ 
from most compounds of other elements. It is convenient, 
therefore, to distinguish these compounds by the term 
organic and to study them under the comprehensive title of 
Organic Chemistry or the Chemistry of Carbon Compounds. 
Several organic compounds have already been discussed in 
the chapter immediately preceding. A few typical com- 
pounds only can be considered in the present chapter. 

Composition of Organic Compounds. — The number of 
organic compounds is very large, but they contain only a few 
elements — seldom more than four or five. Hydrocarbons, 

310 



OTHER CARBON COMPOUNDS 311 

as already stated, contain only carbon and hydrogen. Vege- 
table substances, typified by starch, sugar, and fruit acids, 
contain carbon, hydrogen, and oxygen. Animal substances, 
like hair, albumin, gelatin, and muscle generally contain 
nitrogen as well as carbon, hydrogen, and oxygen; some also 
contain sulphur or phosphorus. Artificial organic com- 
pounds, like dyestuffs and medicines, may contain any ele- 
ment, especially chlorine, iodine, and certain metals. 

The number and complexity of organic compounds are due 
to several facts. (1) Atoms of carbon, unlike those of most 
elements, have power to unite with themselves. (2) Atoms of 
different elements can be introduced into carbon compounds. 
Sometimes these atoms are simply added, sometimes they 
replace other atoms, thus producing an endless number of 
addition and substitution products. (3) The same number 
of atoms may arrange themselves differently, thereby pro- 
ducing isomeric compounds having different properties. 
(4) Organic compounds contain radicals. These radicals 
are groups of atoms analogous to hydroxyl (OH) and ammo- 
nium (NH 4 ), and like these radicals they exist only in com- 
bination. They act chemically like single atoms and pass 
from one compound to another without decomposition. The 
radical C 2 H 5 is called ethyl. It is present in many organic 
compounds, and its presence in ordinary alcohol gives rise 
to the scientific name, ethyl alcohol. Methyl (CH 3 ) is 
another important radical, and phenyl (C 6 H 5 ) is especially 
common in the benzene series of organic compounds. 

Structure of Organic Compounds. — An extensive study 
of the properties of organic compounds has revealed many 
facts about their constitution, i.e. the structure of their mole- 
cules. Relatively little is known about the shape, size, etc., 
of molecules, but much is known about the grouping of atoms 
and of radicals in the molecules. These facts, which are 



312 INORGANIC CHEMISTRY 

ascertained by experiment and are often too complex to be 
expressed briefly, may be represented by suitable formulas. 
Thus, the ordinary or empirical formula of alcohol is C 2 H 6 0. 
But this formula tells nothing about the relation these atoms 
bear to each other. Experiment shows that (1) one hydro- 
gen atom acts differently from the other five, and (2) one 
hydrogen atom is always associated with the oxygen atom in 
chemical changes. Hence, the formula C 2 H 6 . OH expresses 
more fully these facts. Such a formula is called a rational 
or constitutional formula. Sometimes constitution is ex- 
pressed by a graphic or structural formula. Thus, methane 
and ethane have the graphic formulas: — 

H H H 

I .11 

H— C— H H— C-C— H 

I II 

H H H 

Methane Ethane 

In these diagrams the single lines represent a valence of one 
— nothing else, and the number of lines connected with each 
atom is equal to the valence of the element in the compound. 
The lines are sometimes called bonds or links, but they are 
not intended to represent attraction or any other force. The 
graphic formula of ethyl alcohol is : — 

H H 

I I 
H— C— C— O— H 

I I 
H H 

This is not an arbitrary arrangement; the facts mentioned 
above necessitate this general configuration. 

Classification of Organic Compounds. — Organic com- 
pounds are divided and subdivided into many classes, mainly 



OTHER CARBON COMPOUNDS 313 

for the purposes of study. The most common classes are : 
(1) Hydrocarbons ; (2) Alcohols ; (3) Aldehydes ; (4) Ethers ; 
(5) Acids; (6) Esters; (7) Fats, glycerin, and soap; 
(8) Carbohydrates ; (9) Benzene and its derivatives ; 
(10) Cyanogen and its derivates; (11) Proteins. Some 
compounds are closely related and belong to several of 
these groups, while a few common ones are excluded. 

Hydrocarbons 

Three hydrocarbons (methane, ethylene, and acetylene) 
have been considered in Chapter XVI. (See also Benzene, 
below.) 

Three substitution products of methane are chloroform (CHC1 3 ), 
iodoform (CHI 3 ), and carbon tetrachloride (CC1 4 ). Chloroform 
is a heavy liquid. It is made by treating alcohol or acetone with 
bleaching powder, and is used as an anaesthetic. Iodoform is 
a yellow solid. It is made by treating alcohol or acetone with 
iodine and sodium carbonate, and is used as an antiseptic dressing 
for wounds. Carbon tetrachloride is a heavy liquid. It is made 
by passing dry chlorine into carbon disulphide (in which a little 
iodine acts as a catalyzer). It is used to extract fatty substances 
from seeds, bones, and wool ; certain non-inflammable mixtures 
used for cleansing fabrics (e.g. " carbona ") contain carbon tetra- 
chloride. When heated, it forms a heavy, non-inflammable vapor, 
and hence it is used in some fire extinguishers (e.g. " pyrene "). 

Alcohols 

Ordinary or ethyl alcohol is the best-known member of 
this group. It is often called simply alcohol. There are 
many alcohols analogous to ethyl alcohol, but the only other 
important one is methyl alcohol. 

The alcohols may be regarded as hydroxides of certain radicals, 
viz. ethyl, methyl, propyl, etc. 1 For example, ethyl alcohol is ethyl 
hydroxide, and may be considered as formed by replacing one hydro- 

1 The names of these and similar radicals are derived from the name of 
the corresponding hydrocarbon. Thus, methyl from methane, ethyl from 
ethane, propyl from propane. 



314 INORGANIC CHEMISTRY 

gen atom of ethane (C 2 H 6 ) by one hydroxyl group (OH). Again, 
alcohols are analogous to metallic hydroxides, in which the metal 
is replaced by a radical, thus : — 

C 2 H 5 .OH NaOH 

Ethyl Hydroxide Sodium Hydroxide 

Alcohols and metallic hydroxides have some properties in common. 
Thus, both form salts with acids. With acetic acid, sodium hydrox- 
ide forms sodium acetate, while alcohol forms ethyl acetate. (See 
Esters.) 

Methyl Alcohol, CH 3 . OH, is a colorless or slightly yellow- 
ish liquid, much like ordinary alcohol. It boils at about 
66° C, and burns with a pale flame which deposits no soot. 
Methyl alcohol causes blindness and even death. It mixes 
with water in all proportions. It is cheaper than ethyl 
alcohol, and is used as a solvent for fats, oils, and shellac, 
and in the manufacture of varnishes and dyestuffs. Methyl 
alcohol is often called wood alcohol or wood spirit, because 
it is one of the products obtained by the dry distillation of 
wood. (See Charcoal.) 

Ethyl Alcohol, C 2 H 5 . OH, is a colorless, volatile liquid, 
having a burning taste and a pleasant odor. Its specific 
gravity is about 0.8. It boils at about 78° C, and freezes 
at about — 112° C. Alcohol mixes with water in all propor- 
tions. The commercial variety contains about 95 per cent 
of alcohol by volume. Absolute alcohol contains over 99 
per cent of alcohol and is prepared by distilling ordinary 
alcohol with lime. Denatured alcohol is a mixture of 100 
parts ethyl alcohol, 10 parts methyl alcohol, and a small pro- 
portion of benzine or pyridine (or a similar mixture). It 
is unfit for drinking, largely on account of the disagreeable 
taste, but is suitable for industrial uses. Alcohol burns with 
a hot, nearly colorless, non-smoking flame, and is often used as 
a source of heat. It is an excellent solvent for gums, oils, 
and resins, and is therefore extensively used in the manufac- 



OTHER CARBON COMPOUNDS 315 

ture of varnishes, essences, extracts, tinctures, perfumes, 
and medicines. Many organic compounds, as ether and 
chloroform, are prepared from alcohol. Some vinegar is 
made from alcohol. In museums alcohol is used to preserve 
certain specimens. Alcohol is manufactured by the fermen- 
tation of sugars and starches. 

Fermentation is a general term for the chemical changes 
caused by compounds secreted by ferments. The latter are 
minute living organisms ; the compounds they secrete are 
called enzymes. The process and essential products vary 
with the nature of the ferment. The important kinds of 
fermentation are alcoholic, acetic, and lactic ; and they pro- 
duce alcohol, acetic acid, and lactic acid. Alcoholic fermen- 
tation is caused by the enzyme zymase that is secreted by 
ordinary yeast. When yeast is added to a solution of glu- 
cose, maltose, or any other fermentable sugar, the yeast 
plants multiply rapidly. The changes are numerous and 
complex, but the main products resulting from the action of 
the enzyme from the yeast upon the sugar are alcohol and 
carbon dioxide, thus : — 

C 6 H 12 6 = 2C 2 H 6 + 2C0 2 

Dextrose Alcohol Carbon Dioxide 

Commerical alcohol is made from starch. The starch is 
changed into maltose, etc., by an enzyme called diastase, in 
the malt, and the maltose is changed into alcohol and carbon 
dioxide by the zymase in the yeast. Wine, beer, and distilled 
liquors are essentially mixtures of alcohol and water. They 
differ mainly in their proportion of alcohol. The particular 
flavor is due to small quantities of different substances which 
are intentionally added, obtained from the raw materials, or 
formed by special processes of manufacture. Beer contains 
from 3 to 7 per cent of alcohol, wines from 6 to 20, rum, 
brandy, and whisky from 40 to 60 or more per cent. 



316 INORGANIC CHEMISTRY 

Aldehydes and Ketones 

Formaldehyde, CH 2 0, is a gas, but is usually used in solu- 
tion. It has a penetrating odor. The commercial solution 
sold as formalin contains 40 per cent of formaldehyde. 
Formaldehyde is used in the manufacture of dyestuffs and 
fuming nitric acid, and as a disinfectant. When used for 
the last purpose, formalin is vaporized in a special apparatus 
and the vapor is conducted into the infected room. It hard- 
ens tissue and is used as a preservative in museums and 
biological laboratories. Other aldehydes are benzaldehyde 
(oil of bitter almonds, C 7 H 6 0) and vanillin (C 8 H 8 3 ) ; both 
are used as flavors. 

Acetone, C 3 H 6 0, is a colorless liquid which has an ethereal 
odor. It boils at about 56° C. and mixes in all proportions 
with water, alcohol, and ether. It is used as a solvent for 
fats, oils, and waxes, and in the preparation of smokeless 
powders and certain organic compounds. Acetone is one of 
the products obtained by the dry distillation of wood. (See 
Charcoal.) 

Ethers 

Ordinary or ethyl ether is the best-known member of this 
group. 

Ethyl Ether, C 4 H 10 O, is a colorless, volatile liquid, with a 
peculiar, pleasing taste and odor. It boils at 35° C, and the 
vapor is very inflammable. The liquid should never be 
brought near a flame. It is somewhat soluble in water, and 
it also dissolves water to a slight extent. It mixes with 
alcohol in all proportions. It is a good solvent for waxes, 
fats, oils, and other organic compounds. Its chief use is as 
an anesthetic. Ether is manufactured by distilling a mix- 
ture of ethyl alcohol and sulphuric acid in the proper propor^ 
tions, Hence the name, ethyl or sulphuric ether, 



OTHER CARBON COMPOUNDS 317 

Acids 

This large class of compounds is divided into several series, 
one of the most important of which is the acetic or fatty 
series. Its best known member is acetic acid; several of the 
higher members occur in fats and oils. These acids are 
closely related to hydrocarbons, alcohols, and aldehydes, as 
may be seen by the following formulas : — 

H H H 



H— C— H 


H— C— (OH) 


c=o 


0=C— (OH) 


H— C— H 


H— C— H 

1 


H— C— H 

i 


H— C— H 

I 


i 
H 


1 

H 


1 
H 


1 
H 


Ethane 


Ethyl Alcohol 


Acetic Aldehyde 


Acetic Acid 



The characteristic radical of organic acids is COOH (or 
0=C — O — H), and is called carboxyl. 
I 

Acetic Acid, C 2 H 4 2 or CH 3 . COOH. This is the most 
common organic acid. It is manufactured on a large scale 
by the dry distillation of wood. The dark red watery dis- 
tillate, which is called pyroligneous acid, contains about 10 
per cent of acetic acid, besides methyl alcohol and acetone. 
This distillate is neutralized with lime or sodium carbonate, 
and the acetate formed is then decomposed and distilled with 
sulphuric acid. The acetic acid which condenses in the re- 
ceiver may be further purified by distilling it with potassium 
dichromate and then filtering through charcoal. Sometimes 
the pyroligneous acid is distilled without neutralizing; the 
distillate is then dilute, impure acetic acid, and is known as 
wood vinegar. If sodium acetate, prepared as described 
above, is fused and then distilled with concentrated sul- 
phuric acid, the product is very concentrated acetic acid. 



318 INORGANIC CHEMISTRY 

It is called glacial acetic acid, because at about 17° C. it 
becomes an icelike solid. 

Commercial acetic acid is a water solution containing 
about 30 per cent of pure acetic acid. It is a colorless liquid, 
having a pleasant odor and a sharp taste. It is a weak acid, 
a normal solution at 18° C. being dissociated to the extent 
of about .4 per cent into the ions H + and C 2 H 3 2 ". It 
mixes with water and alcohol in all proportions, and like 
alcohol is an excellent solvent for many organic substances. 
Recently, it has begun to replace alcohol as a solvent for 
many drugs. 

Acetic acid is used to prepare acetates, dyestuffs, medi- 
cines, white lead, and in the manufacture of vinegar. 

Vinegar is dilute, impure acetic acid. It is prepared by oxidizing 
dilute alcohol, the essential change being represented thus : — 

C 2 H 6 + 2 = C2H4O2 + H 2 

Alcohol Oxygen Acetic Water 

Acid 

The transformation is accomplished by fermentation. (1) When 
beer, weak wines, or cider are exposed to the air, they slowly become 
sour, owing to the conversion of alcohol into acetic acid. The change 
is caused by the presence and activity of a ferment, known as myco- 
derma aceti, or "mother of vinegar." Strong wines and pure dilute 
alcohol do not become sour, because the ferment cannot live in 
such liquids. (2) Fruit juices . and molasses contain fermentable 
sugar and ferment when exposed to the air (which always contains 
the necessary organisms), forming alcohol first and finally vinegar. 
Cider vinegar is made this way. (3) In the " quick vinegar pro- 
cess," impure dilute alcohol is oxidized to acetic acid by exposing it 
to an excess of air. The operation is conducted in tall vats or casks 
filled with beechwood shavings soaked in old vinegar. Holes at the 
bottom and top allow air to enter and escape freely. The alcoholic 
solution is introduced at the top, trickles through the shavings, and 
collects at the bottom. In its passage it comes in contact with the 
ferment and oxygen, and is partially converted into vinegar. The 
operation is repeated until the change is complete.- Thus prepared, 
the vinegar lacks the flavor, odor, and color of cider vinegar, but these 
deficiencies may be artificially supplied. 



OTHER CARBON COMPOUNDS 319 

Acetates. — Acetic acid is a monobasic acid, and forms only 
one series of salts — the acetates. They are prepared, like 
other salts, by the interaction of the acid and carbonates, 
hydroxides, metals, etc. The metallic acetates are usually 
crystalline solids, which readily yield acetic acid when 
treated with sulphuric acid.. Most of them contain water of 
crystallization, and several are poisonous. 

Acetates have many applications. Sodium acetate, NaC2Hs02. 3 H2O, 
is a white crystalline solid, used in preparing pure aectic acid and in 
the manufacture of dyestuffs. Lead acetate, Pb(C2H302)2, is a white 
crystalline solid, used in dyeing and in making a yellow pigment. Its 
sweet taste led to the common name of " sugar of lead." Aluminium 
acetate, A1(C2H 3 2 )3, is not known in the pure state, but an impure 
solution, known as " red liquor," is extensively used in dyeing and in 
calico printing. Iron acetates are sold in solution as a complex black 
liquid, known as " iron liquor," which is used in dyeing black silks 
and cottons, and in calico printing. (See Mordants.) A complex 
copper acetate, 2 Cu(C2H 3 02)2+CuO, called verdigris, is used in making 
blue paint. Another complex acetate of copper and arsenic is Paris 
green ; it is used to kill potato bugs and other injurious insects. 

Other Organic Acids are oxalic, lactic, malic, butyric, stearic, 
palmitic, oleic, tartaric, and citric. 

Oxalic Acid, C2H2O4, occurs as a calcium salt (CaC 2 4 ) in many 
plants, e.g. rhubarb. Oxalic acid is a white solid, very soluble in 
water, from which it crystallizes with two molecules of water of 
crystallization (C2H2O4 . 2 H 2 0). It is very poisonous. The acid 
and some of its salts decompose iron rust and inks containing iron, 
and are used to remove stains. 

Lactic Acid, C 3 H 6 3 , occurs in sour milk (see page 326). 

Malic Acid, C 4 H 6 5 , occurs free or as salts in many fruits and 
in parts of vegetables. 

Butyric Acid, C 4 H 8 2 , occurs in rancid butter. Stearic Acid, 
Ci 8 H 3 60 2 , and Palmitic Acid, C16H32O2, are white solids. Oleic 
Acid, C18H34O2, is an oily liquid. Derivatives of these four acids 
occur in fats and oils (see page 322). 

Tartaric Acid, C 4 H 6 6 , occurs as the acid potassium salt 
(HKC 4 H 4 6 ) in grapes and other fruits. During the fermentation 
of grape juice, impure acid potassium tartrate is deposited in the 
casks. From this argol or crude tartar the acid itself is prepared. 



320 INORGANIC CHEMISTRY 

Tartaric acid is a white crystalline solid, soluble in water and 
alcohol. It is used in dyeing, and as one ingredient of Seidlitz 
powders. In these and similar mixtures it serves to decompose the 
other ingredient, which is a carbonate. (See Sodium Bicarbonate.) 

Tartaric acid is dibasic and forms two classes of salts. Purified 
acid potassium tartrate obtained from argol is commonly known as 
cream of tartar (HKC4H4O6). It is extensively used in the manu- 
facture of baking powders. These, as a rule, are essentially mix- 
tures of cream of tartar, sodium bicarbonate (HNaCOs), and a little 
starch. When moistened by dough, the baking powder dissolves, 
the acid salt and the carbonate interact and liberate carbon dioxide. 
This gas escapes slowly through the dough, thereby purring it up 
and making it porous. (See Sodium Bicarbonate.) Tartar emetic is 
potassium antimonyl tartrate (KSbOC4H40 6 ). It is used as a medi- 
cine and to some extent in dyeing. Rochelle salt is potassium 
sodium tartrate (KNaC4H40e). 

Citric Acid, C 6 H807, occurs abundantly in lemons and oranges, 
and in small quantities in currants, gooseberries, and raspberries. It 
is a white crystalline solid, very soluble in water. The taste is sour, 
but pleasant. The acid and its magnesium salt are used as medi- 
cines. The acid itself is used in calico printing. Citric acid is tribasic. 

Esters 

Esters are compounds of carbon, hydrogen, and oxygen 
closely related to alcohols and organic acids. Thus, when 
ethyl alcohol, acetic acid, and concentrated sulphuric acid 
are mixed and warmed, ethyl acetate is one product. The 
essential change is represented thus : — 

C 2 H 5 . OH + CH 3 . COOH = CH 3 . COOC 2 H 5 + H 2 

Ethyl Alcohol Acetic Acid Ethyl Acetate Water 

Ethyl acetate has a pleasant, fruitlike odor, and its forma- 
tion in this way is a simple test for alcohol or acetic acid. 
Ethyl acetate is analogous to sodium acetate, i.e. the organic 
salt contains the radical ethyl, while the metallic salt contains 
sodium. The fatty acids, as well as those of other series, 
form many esters of special interest. Some occur naturally 



OTHER CARBON COMPOUNDS 321 

in fruits and flowers, and in many cases give the fragrance 
and flavor. Others are prepared artificially and used as the 
characteristic ingredient of cheap flavoring extracts, per- 
fumery, and beverages. Ethyl butyrate has the taste and 
fragrance of pineapples, amyl acetate of bananas, amyl 
valerate of apples, methyl salicylate of wintergreen. 

Fats, Glycerin, and Soap 

General Relations. — • Natural fats and oils are essentially 
mixtures of stearin, palmitin, and olein. Beef and mutton 
fat are chiefly stearin, lard is mainly palmitin and olein, while 
oils such as olive oil are largely olein. Stearin and palmitin 
are solids at the ordinary temperature, but olein is a liquid. 
These three compounds — stearin, palmitin, and olein — 
are esters of their corresponding acids and the alcohol gly- 
cerin. They are analogous to ethyl acetate. The radical 
of glycerin is glyceryl, C 3 H 5 . Stearin is glyceryl stearate, 
palmitin is glyceryl palmitate, and olein is glyceryl oleate. 
Natural fats and oils, therefore, are mixtures of these and 
similar esters. Glycerin is a triacid alcohol containing three 
hydroxyl (OH) groups. Like ordinary alcohol, it interacts 
with the fatty acids and forms esters. The latter, as we have 
just seen, are the fats. Now, when fats are heated with very 
hot steam or with sulphuric acid, they -are changed into 
glycerin and the corresponding acids. Thus, with stearin 
the change is — 

(C 17 H 35 . COO) 3 C 3 H 5 + 3 H 2 = C 3 H 5 (OH) 3 + 3 C 17 H 35 . COOH 

Stearin Glycerin Stearic Acid 

But if fats are boiled with sodium hydroxide or a similar 
alkali, glycerin and an alkaline salt of the corresponding 
acid are formed. Soap is a mixture of such alkaline salts. 
In a few words, the general relations are these : (1) fats are 
esters; (2) treated with steam or acid, fats form glycerin 



322 INORGANIC CHEMISTRY 

and fatty acids; (3) treated with alkalies, fats form glycer- 
in and soap. 

Natural Fats and Oils are often complex mixtures. The 
solid fats, as already stated, are rich in stearin and palmitin. 
Tallow is chiefly stearin, but human fat and palm oil are 
largely palmitin. The soft and liquid fats and oils contain 
considerable olein, as a rule. The proportion of olein deter- 
mines the consistency of the fats and oils. Thus, olive oil 
contains 72 per cent of olein (and a similar fat) and about 
28 per cent of stearin and palmitin. The specific character 
of many fats and oils is due mainly to a small proportion of 
certain fats. These fats correspond to uncommon acids in 
the fatty, oleic, and other series. Butter, for example, con- 
sists of the fats corresponding to the following acids : palmitic, 
stearic, oleic, butyric, capric, and caproic. The last three, 
together with traces of other substances, give butter its 
pleasant flavor. Oleomargarine and other substitutes for 
butter resemble real butter very closely in composition. 
Artificial butter, however, lacks the flavor of the real butter, 
but it is " probably just as nutritious, although perhaps 
not quite so easily digested. " 

Glycerin, C 3 H 8 3 or C 3 H 5t . (0H) 3 , is a thick, sweet liquid. 
It mixes readily with water and with alcohol in all propor- 
tions, and absorbs moisture from the air. Heated in the air 
it decomposes and gives off irritating gases, like those pro- 
duced by burning fat. 

Glycerin is used to make nitroglycerin (see below), toilet 
soaps, and printer's ink rolls ; it is also used as a solvent, 
a lubricator, a preservative for tobacco, and certain foods, 
a sweetening substance in certain liquors, preserves, and 
candy ; as a cosmetic ; and owing to its non-volatile and 
non-drying properties, it is used as an ingredient of certain 
inks and oils. 



OTHER CARBON COMPOUNDS 323 

Glycerin is a by-product in the manufacture of soap, or it is 
made directly by decomposing fats with steam under pressure or 
with lime. All these methods involve the chemical change described 
above, viz. the decomposition of an ester (the fat) into the cor- 
responding alcohol (glycerin) and a mixture of fatty acids. By 
skillful treatment the glycerin is freed from the water and im- 
purities. 

As already stated, glycerin is an alcohol, and for this reason it 
is often called glycerol. When treated with a mixture of concen- 
trated nitric and sulphuric acids, it forms an ester commonly known 
as nitroglycerin (C 3 H 5 (ON0 2 )3). This is a yellow, heavy, oily 
liquid. It is the well-known explosive, and is also an ingredient of 
some other explosives. When kindled by a flame, it burns with- 
out explosion ; but if it is compressed, detonated, or heated to 
about 250° C, it explodes violently. Nitroglycerin is dangerous to 
handle and transport, and is usually mixed with some porous sub- 
stance, such as infusorial earth. In this form it is called dynamite. 
Other explosives contain nitroglycerin, e.g. blasting gelatin and 
cordite. 

Soap, as already stated, is a mixture of alkaline salts of 
organic acids, mainly stearic and palmitic acids. Soap is 
made by boiling fats with sodium hydroxide or potassium 
hydroxide. This process is called saponification. Sodium 
hydroxide produces hard soap, consisting chiefly of sodium 
palmitate, sodium stearate, and sodium oleate. Potassium 
hydroxide produces a soft, semi-fluid soap, which contains 
mainly the corresponding potassium salts. The chemical 
change, as already stated, consists in the transformation of 
an ester (fat) into glycerin and an alkaline salt. In the case 
of pure stearin (glyceryl stearate) the change may be repre- 
sented thus : — 

CsH5(Ci7H8B.COO)8+3NaOH=3Ci7H35COONa+C3H B (OH)8 

Stearin Sodium Sodium Glycerin 

Hydroxide Stearate 

The fats used in soap making vary. Tallow, lard, palm oil, 
and cocoanut oil make white soaps. Grease, together with 



324 INORGANIC CHEMISTRY 

tallow, palm oil, cottonseed oil, and rosin, make yellow soaps. 
Olive oil is used for making castile soap. 

Most soaps are manufactured by the boiling process. The fat 
and alkali are boiled in a huge kettle. This operation produces a 
thick, frothy mixture of soap, glycerin, and alkali. At the proper 
time salt is added, thereby causing the soap to separate and rise to 
the top. The liquid beneath is drawn off, and from it glycerin is 
extracted. The soap is often boiled again with rosin or cocoanut oil ; 
then mixed, if desired, with perfume, coloring matter, or some 
filling material (such as sodium silicate, sand, or borax). Floating 
soaps are made by forcing air into the semi-solid soap before cool- 
ing. The best soaps do not contain unchanged fat or " free alkali," 
i.e. sodium hydroxide. 

The cleansing action of soap is probably due to two causes. 
(1) Soap hydrolyzes with water and the liberated sodium hydroxide 
acts upon the grease and oil that are mixed with the dirt. (2) Soap 
causes oils to form an emulsion which is readily removed by water. 
Doubtless the second cause is the more efficient. 

Carbohydrates 

Sugar. — The popular term sugar means almost any 
sweet substance found in fruits, nuts, vegetables, sap of 
trees, etc., though it is usually restricted to the ordinary 
white sugar obtained from sugar cane and sugar beet. 
Chemically, there are many sugars, each having a definite 
constitution. The most important is ordinary sugar, which 
is also called cane sugar, sucrose, and saccharose. Other 
important sugars are dextrose, levulose, lactose, and maltose. 

Cane Sugar, C12H22O11, is widely distributed in nature, 
being found in the sugar cane, sugar beet, sugar maple, 
Indian corn, sorghum, most sweet fruits, many nuts, blos- 
soms of flowers, and honey. The main source of cane sugar 
is the sugar cane and sugar beet. 

Cane sugar is a white, crystalline solid. It is very soluble 
in water, one part of water dissolving about three times its 



OTHER CARBON COMPOUNDS 325 

weight of sugar at ordinary temperatures. If heated to about 
160° C, it melts, and on cooling becomes a glassy solid. As 
the temperature is raised, the solid begins to decompose, 
and at about 210° C. water is given off and a brown substance 
called caramel is formed. Further heating produces a black 
porous mass of carbon called sugar charcoal. 

The manufacture of sugar from sugar cane and sugar beets 
involves two main operations. (1) In the preparation of raw sugar 
from sugar cane the juice obtained by crushing the cane is first boiled 
with a weak calcium hydroxide solution to neutralize acids, remove 
impurities, and prevent fermentation, next freed from excess of lime 
by carbon dioxide, and finally filtered through bone black. The 
purified juice is then evaporated in vacuum pans until the sugar be- 
gins to crystallize from the cooled liquid. The crystals are then 
separated from the brown liquid by a centrifugal machine. The 
liquid is the familiar molasses. In the preparation of raw sugar 
from sugar beets the washed beets are cut into slices and soaked in 
water to dissolve the sugar. The solution is treated much like cane 
sugar solutions. (2) Raw sugar is dark colored, and must be re- 
fined before it is suitable for most uses. The raw sugar is first dis- 
solved in water, and lime and other substances are added to gather 
the impurities into a scum or clot. The colored liquid is next 
filtered, first through cloth bags and then through bone black. 
The filtered sirup is evaporated in large vacuum pans until a sample 
deposits the right size crystals. The crystals of sugar are separated 
from the sirup by centrifugal machines, then dried and separated 
in a heated tube called a granulator. Hence the name granulated 
sugar. 

Dextrose and Levulose. — When sucrose is heated with 
dilute acids, the two sugars dextrose and levulose are formed. 
The chemical change is an example of hydrolysis and may 
be represented thus : — 

C12H22O11 + H2O = CeHi 2 06 + CeH^Oe 
Sucrose Dextrose Levulose 

The same change is brought about by an enzyme called 
invertase. Dextrose is a white solid about three fifths as 



326 INORGANIC CHEMISTRY 

sweet as sucrose. Dextrose is found in honey and in many 
fruits, especially grapes, and is sometimes called grape sugar. 
Another name for it is glucose. Levulose is also a sweet, 
white solid found in fruits and honey, and is often associated 
with dextrose. It is sometimes called fructose or fruit sugar. 

Commercial glucose contains about 40 to 50 per cent of dextrose. 
It is manufactured by heating starch with dilute sulphuric acid. 
The starch is first changed into a sweet solid called dextrin, then into 
dextrose, and if the process is carried far enough, the product is a 
hard, waxlike solid known as commercial grape sugar, which is 
almost pure dextrose. Glucose is an inexpensive substitute for 
sucrose and is extensively used in making candy, jellies, sirups, and 
other sweet mixtures. 

Dextrose, and also levulose, is converted by yeast into ethyl 
alcohol and carbon dioxide (see Alcohol). Dextrose and levulose 
are reducing agents. An alkaline solution of dextrose is used to 
reduce a silver solution and deposit the silver as a bright film in 
making reflectors, mirrors, Dewar flasks, and thermos bottles. 
It also reduces a strongly alkaline mixture of copper sulphate and 
sodium potassium tartrate, known as Fehling's solution. When this 
solution is boiled with dextrose (or any other reducing sugar), a 
reddish copper compound (cuprous oxide, Cu 2 0) is formed. This 
experiment is often used as a test for dextrose and similar sugars. 
Solutions of dextrose and levulose rotate the plane of polarized 
light — dextrose to the right and levulose to the left. That is, 
when their solutions are placed in a sugar-polariscope and examined, 
the light instead of passing entirely through the instrument is 
extinguished ; and in order to bring about illumination again, the 
plane of the polarized light must be rotated a certain number of 
degrees in order to compensate for the rotation caused by the sugar 
solution. By means of this instrument valuable information can 
be obtained about the kind and proportion of sugar in solutions. 

Lactose (milk sugar, Ci 2 H 22 0ii . H 2 0) occurs in the milk of 
mammals. Cow's milk contains on the average 4.88 per cent of 
lactose. Lactose is not so sweet or soluble as cane sugar. When 
milk sours, its lactose changes into alcohol and lactic acid. The 
acid causes the sour taste and also assists in curdling the milk, i.e. 
in changing the casein into a clot or curd. Casein is used in making 
cheese. 



OTHER CARBON COMPOUNDS 327 

Maltose is formed from starch by malt, hence the name maltose. 
The transformation is caused by the enzyme diastase. Maltose 
ferments readily with yeast, forming alcohol and carbon dioxide, 
and is manufactured in large quantities for the commercial produc- 
tion of alcohol and fermented liquors. With dilute acids, maltose 
forms dextrose by hydrolysis. Like lactose, maltose is a sweet 
solid, very soluble in water, from which it forms crystals (C12H22O11 . 
H 2 0) ; its solution turns the plane of polarized light to the right and 
reduces Fehling's solution. 

Starch is widely distributed in the vegetable kingdom. 
It is found in wheat, corn, and all other grains ; in potatoes, 
beans, peas, and similar vegetables ; and in large quantities 
in rice, sago, tapioca, and nuts. Many parts of plants con- 
tain starch. 

Starch is a white mass, as usually seen. But under the 
microscope it is found to consist of oval grains varying some- 




f9 

Fig. 50. — Starch grains (magnified) — wheat (left) , rice (center) , corn (right). 

what with the source (Fig. 50). Starch is only very slightly 
soluble in water. But if heated with water, the grains swell 
and burst, partially dissolve, and form a solution which, when 
cold, becomes the familiar starch paste. Starch in solution 
is turned blue by iodine, and its presence in many vegetables 
and foods can be readily shown by grinding the substance in 
a mortar with cold water and adding a drop of dilute iodine 
solution. The composition of starch corresponds to the 
formula (C 8 Hio0 5 ) H . 



328 INORGANIC CHEMISTRY 

Bread. — Wheat flour contains about 70 per cent of starch. 
The remainder is chiefly water and gluten, though small 
quantities of mineral matter and fat are present. In making 
bread, the flour, water, and yeast are thoroughly mixed into 
dough, which is put in a warm place to rise. Fermentation 
begins at once. The enzymes change the starch into dextrose 
or a similar fermentable substance which undergoes fer- 
mentation, forming alcohol and carbon dioxide. The gases 
escape in part through the dough, which becomes light and 
porous. When the dough is baked, the heat kills the yeast 
plant and fermentation stops ; but the alcohol, carbon diox- 
ide, and some water escape and puff up the mass still more. 
The heat, however, soon hardens the starch, gluten, etc., 
into a firm but porous loaf. 

Cellulose, (C 6 Hi O5)„, is the basis of the cells of plants. 
Wood, cotton, linen, and paper are largely cellulose. Pure 
cellulose is a white substance, insoluble in most liquids, but 
soluble in a mixture of ammonia and copper oxide. Concen- 
trated sulphuric acid dissolves it slowly ; and if the solution 
is diluted and boiled, the cellulose is changed into a mixture 
of glucose and dextrin. Sulphuric acid of a special strength, 
if quickly and properly applied to paper, changes it into a 
tougher form called parchment paper. Ordinary paper 
consists chiefly of cellulose matted together. 

Cellulose with nitric acid forms cellulose nitrates. One of the 
cellulose nitrates is gun cotton. It looks like ordinary cotton, and 
may be spun, woven, and pressed into cakes. It burns quickly, 
if unconfined ; but when ignited by a percussion cap or when burned 
in a confined space, gun cotton explodes violently. It is used in 
blasting and for torpedoes and submarine mines. A mixture of 
gun cotton, alcohol, and ether forms a transparent solid called smoke- 
less powder ; when exploded it yields colorless gases. A solution 
of certain cellulose nitrates in a mixture of alcohol and ether is 
called collodion. It is used in preparing certain photographic 



OTHER CARBON COMPOUNDS 329 

material and as a coating for wounds. A mixture of camphor and 
cellulose nitrates is called celluloid. 

Benzene and its Derivates 

Benzene, C 6 H 6 , is a colorless liquid, lighter than water 
(sp. gr. .88 at 20° C.) and boils at 80° C. It burns with a 
luminous, smoky flame, owing to its richness in carbon. 
Ordinary illuminating gas owes its luminosity partly to ben- 
zene. It dissolves fats, resins, iodine, sulphur, and rubber. 
Benzene is sometimes called benzol. It should not be con- 
fused with benzine, w T hich is a mixture of hydrocarbons de- 
rived from petroleum. Benzene is chiefly used in preparing 
its derivatives. 

Nitrobenzene, C 6 H 5 . N0 2 , is a yellow liquid formed by the inter- 
action of benzene and nitric acid. The equation for the chemical 
change is : — 

C 6 H 6 + HNOa = CsHs.NOj + H 2 

Benzene Nitric Nitro- Water 

Acid benzene 

It is volatile, and the vapor, which is poisonous, has the odor of 
bitter almonds. It is chiefly used in the manufacture of aniline. 

Aniline, C 6 H 5 . NH 2 , is an oily liquid, slightly heavier than water. 
It is prepared on a large scale by reducing nitrobenzene with nascent 
hydrogen, thus : — 

C 6 H 5 . N0 2 + 6 H = C 6 H 5 . NH 2 + 2 H 2 

Nitrobenzene Hydrogen Aniline Water 

From aniline are made many compounds known as aniline dyes. 

Phenol, C 6 H 5 . OH, is a white crystalline solid. It has a smoky 
odor, is poisonous, and burns the skin. Coal tar is the source of 
phenol. A solution of phenol in water, popularly called carbolic 
acid, is used as a disinfectant. 

Naphthalene, Ci H 8 , is a white, lustrous, crystalline solid ob- 
tained from coal tar. It has a penetrating, unpleasant odor, and is 
used as a substitute for camphor under the name* of " moth balls." 
Large quantities of naphthalene are used in making indigo. 



330 INORGANIC CHEMISTRY 

Anthracene, Ci 4 Hi , is a white crystalline solid, and is obtained 
from coal tar. It is one of the most important hydrocarbons, be- 
cause from it alizarin is made. Alizarin is a valuable dyestufx", 
because it produces brilliant, fast colors with different mordants. 

Cyanogen and its Derivatives 

Cyanogen, (CN) 2 , is a colorless gas, has the odor of peach 
kernels, is poisonous, and burns with a purplish flame. 

Hydrocyanic or prussic acid, HCN, is prepared by heating 
a cyanide with sulphuric acid. The acid smells like peach 
kernels, and is extremely poisonous. It is a feeble acid and 
dissociates slightly into H + and CN". 

Potassium cyanide, KCN, and sodium cyanide, NaCN, 
are white, deliquescent solids. They are very poisonous. 
Large quantities are used in gold and silver plating and in 
the cyanide process of extracting gold from its ores, as de- 
scribed under that metal. The alkaline reaction exhibited 
by a solution of potassium cyanide is due to hydrolysis. 

Potassium sulphocyanate or thiocyanate, KCNS, is a white 
crystalline salt, which produces a deep red solution (due to 
ferric sulphocyanate) with solutions containing ferric ion. 
(See Ferric Compounds.) 

Proteins 

Proteins are complex nitrogenous compounds often desig- 
nated by the term protein (formerly proteid) . Besides nitro- 
gen, they contain carbon, hydrogen, and oxygen, usually 
sulphur, often phosphorus, and occasionally iron. Common 
proteins contain 15 to 18 per cent of nitrogen. They burn 
with a disagreeable odor arid liberate ammonia. When 
proteins putrefy, poisonous substances called ptomaines are 
often produced. 

Proteins constitute the principal part of the tissue of the 



OTHER CARBON COMPOUNDS 331 

cells of our bodies. The body of an average man is about 18 
per cent protein. Protein in some form must be a part of 
the food of animals, since protein serves to replace worn- 
out tissue and contribute new tissue for the growth, of the 
body. Important groups of proteins are albumins, globulins, 
glutelins and protamins, albuminoids, phosphoproteins, and 
hemoglobins. Members of these groups are widely distrib- 
uted, being found in eggs, milk, seeds of plants (e.g. beans 
and peas), cereals (e.g. wheat, rye, and barley), muscle, 
blood, and animal tissues. 

Problems 

1. Calculate the affinity constant of a molar solution of acetic 
acid in which the proportion ionized is .0041. Ans. .0000169. 

2. What weight of potassium hydroxide will neutralize (a) 100 
gm. of acetic acid and (6) 10 gm. of oxalic acid (dibasic) ? 



CHAPTER XVIII 
Sulphur and its Compounds 

Sulphur has been known for ages. The alchemists re- 
garded it as one of the primary forms of matter. The ele- 
ment and its compounds have always played an important 
part in the development of many industries. 

Occurrence. — Sulphur, free and combined, is abundant 
and widely distributed. Free or native sulphur is found 
usually in volcanic regions. There are also beds associated 
with gypsum (calcium sulphate). It is believed that some 
deposits were formed by the reduction of gypsum by micro- 
organisms. 

Several important metallic ores are native sulphides, e.g. 
lead sulphide (PbS), zinc sulphide (ZnS), and those of mer- 
cury, antimony, and copper. Iron sulphide (FeS 2 ) is also 
plentiful. The most abundant sulphates are calcium sul- 
phate (CaS0 4 ), barium sulphate (BaS0 4 ), and magnesium 
sulphate (MgS0 4 ). Volcanic gases often contain sulphur 
dioxide (S0 2 ) and hydrogen sulphide (H 2 S). The latter is 
also found in sulphur springs. Sulphur is a constituent of 
protein, and hence is present in many kinds of animal and 
vegetable matter, e.g. eggs and mustard. Some varieties of 
petroleum and coal contain sulphur compounds. 

Source. — Until about 1903 Sicily furnished most of the 
sulphur, but the larger part of the world's supply now comes 
from Japan and the United States (especially Louisiana). 
Some of the sulphur of commerce is obtained by roasting 

332 



SULPHUR AND ITS COMPOUNDS 



333 



iron pyrites (FeS 2 ) as in the manufacture of sulphuric acid. 
Small amounts are recovered from the calcium sulphide 
waste of the Leblanc soda process (see Sodium Carbonate), 
and from the residues of the iron oxide used to purify illu- 
minating gas. 



Extraction. — For many years 
sulphur has been extracted from 
the impure native sulphur in Sicily 
by a primitive process. The sul- 
phur ore is brought to the surface 
by laborers, piled loosely in a heap, 
and covered with powdered or 
burnt ore or with earth. The 
heap is ignited at the bottom, and 
the heat produced by the combus- 
tion of some of the sulphur melts 
the rest, which runs out at the 
bottom. In Louisiana, which now 
furnishes most of the sulphur used 




Fig. 51. — Section of a set of 
pipes for winning sulphur. 
The water enters through the 
spaces A and B and melts the 
sulphur, which rises part way 
in C and is forced to the sur- 
face by compressed air intro- 
duced through the innermost 
pipe. 



in the United States, the deposits 
are about half a mile in diameter and 500 feet thick and 
are located about 800 feet beneath the surface. They are 
reached by wells consisting essentially of a set of four con- 
centric pipes (Fig. 51) sunk through a mass of clay, quick- 
sand, and rock. Water heated (under pressure) to about 
170° C. is pumped down the two outer pipes (A, B). After 
some of the sulphur is melted, compressed air is forced down 
through the innermost pipe into the molten mass, thereby 
forcing the mixture of air and melted sulphur up through 
the pipe C. The liquid flows into huge wooden bins where 
it solidifies. The wells are very powerful, a single well 
often pumping 500 tons of sulphur daily. The annual 
production is over 200,000 tons. 



334 



INORGANIC CHEMISTRY 



Purification. — Louisiana sulphur is about 99 per cent 
pure. Crude Sicilian sulphur requires purification. This 
is accomplished by the apparatus shown in Figure 52. The 
crude sulphur is melted in B, and flows into the iron cylin- 




Fig. 52. — Apparatus for purifying sulphur. 

der A. Here it is heated, and the vapors pass into the 
large brick chamber, provided with a tap C, from which 
the liquid sulphur may be withdrawn. If the distillation 
is conducted slowly, the sulphur vapor condenses upon the 
cold walls of the chamber as a fine powder, called flowers 
of sulphur. As the operation continues the walls become 
hot, and the sulphur collects on the floor as a liquid which is 
drawn off into cylindrical wooden molds, forming roll sul- 
phur or brimstone. 

Properties. — Ordinary sulphur is a yellow, brittle, crystal- 
line solid. It is practically insoluble in water, but most 



SULPHUR AND ITS COMPOUNDS 335 

varieties dissolve in carbon disulphide, and to some extent in 
turpentine, chloroform, and benzene (C 6 H 6 ). Sulphur does 
not conduct heat well, the warmth of the hand even causing 
it to crackle and break from the unequal expansion. The 
specific gravity of the solid is about two. The density of 
the vapor varies with the temperature. At the lowest tem- 
perature at which sulphur can be vaporized, one molecule 
contains eight atoms (S 8 ), while at 800° C. it contains two 
atoms (S 2 ). Dissolved sulphur has the formula S 8 . 

Heated to 114.5° C. sulphur melts to a thin pale yellow 
liquid. At about 160° C. the liquid becomes dark brown 
and viscous, and at about 230° C. it is black and too thick 
to be poured from the vessel. Heated still higher, the color 
remains black but the mass becomes thin, and finally at 
about 445° C. the liquid boils and turns into yellow sulphur 
vapor. When cooled slowly, the sulphur undergoes the same 
series of changes in the reverse order. Sulphur ignites 
readily and burns with a pale blue flame, forming sulphur 
dioxide gas (S0 2 ) ; if burned in oxygen, a little sulphur 
trioxide (S0 3 ) is also formed. Finely divided sulphur oxidizes 
in moist air, forming sulphuric acid (H 2 S0 4 ). It combines 
directly and readily with hydrogen, carbon, and chlorine, 
forming hydrogen sulphide, carbon disulphide, and sulphur 
chlorides (S 2 C1 2 and SC1 4 ). The reaction between sul- 
phur and metals is often attended by vivid combustion, 
though heat is necessary to start the chemical action. Thus, 
when a mixture of flowers of sulphur and powdered iron is 
heated, the mass begins to glow and soon becomes red-hot, 
the glow often spreading through the mass after removal 
from the flame. The product is iron sulphide, and the chemi- 
cal change is represented thus : — 



Fe 


+ s 


FeS 


Iron 


Sulphur 


Iron 
Sulphide 



336 INORGANIC CHEMISTRY 

Heated copper glows when dropped into melted sulphur, 
while zinc dust and flowers of sulphur combine violently. 

Different Forms of Sulphur. — Sulphur exists in several 
different forms, which are crystallized or amorphous. These 
modifications differ in specific gravity, solubility, and other 
properties. The crystallized forms belong to the ortho- 
rhombic and monoclinic systems. (See Appendix, § 3.) 
Orthorhombic sulphur is the form obtained by crystalliza- 
tion from a solution of carbon disulphide. Crystallized 
native sulphur and ordinary roll sulphur are orthorhombic, 
though the latter usually consists of such a mass of inter- 
laced crystals that the form is obscured. The monoclinic 
sulphur is the form* obtained by slowly cooling molten sul- 
phur. By melting sulphur in a crucible and pouring off the 
excess of liquid as soon as crystals shoot out from the walls 
near the surface, the interior of the crucible is found to be 
lined with long, dark yellow, shining needles. They are 
monoclinic crystals of sulphur. After a few days they gradu- 
ally become opaque and slowly change into minute ortho- 
rhombic crystals. This change is due to the fact that 
monoclinic sulphur is in the stable form only between 96° 
C. and 119° C. (its melting point), while orthorhombic sul- 
phur is in the stable form only below 96° C. Above 96° C. 
orthorhombic slowly becomes monoclinic, below 96° C. mono- 
clinic slowly becomes orthorhombic. This temperature at 
which the two forms of sulphur pass into one another is 
called the transition point. These two varieties of crystal- 
lized sulphur have different properties. Orthorhombic sul- 
phur has the specific gravity 2.06 and melts at 114.5° C. 
(if heated rapidly) ; the corresponding values of monoclinic 
sulphur are 1.96 and 119.25° C. Amorphous sulphur is 
formed by pouring boiling sulphur into cold water. It is a 
tough, plastic, rubberlike, amber-colored mass, mostly in- 



SULPHUR AND ITS COMPOUNDS 337 

soluble in carbon disulphide. Its formation is due to the 
sudden cooling of the viscous liquid sulphur. It soon be- 
comes hard and yellow, part crystalline and part (about 30 
per cent) amorphous. Amorphous sulphur is sometimes 
found in flowers of sulphur and can be detected by its in- 
solubility in carbon disulphide. Other varieties of sulphur 
also contain amorphous sulphur. One is a white or whitish 
powder, made by boiling flowers of sulphur with milk of 
lime and adding hydrochloric acid to the decanted liquid ; 
a fine sulphur powder is precipitated, which gives the liquid 
the appearance of milk, hence the name often applied to it, 
" milk of sulphur/ J 

Uses. — Sulphur is used in making sulphuric acid and 
other sulphur compounds, gunpowder, fireworks, matches, 
in vulcanizing rubber, and as a medicine. Considerable is 
used to kill phylloxera (an insect which destroys grapevines) ; 
some insecticides, made from sulphur, e.g. lime-sulphur 
sprays, contain unstable compounds which by decomposi- 
tion liberate sulphur upon the insect pest. 

The Important Compounds of Sulphur are hydrogen sul- 
phide and other sulphides (especially metallic sulphides), 
sulphur dioxide and trioxide, the sulphites, sulphuric acid and 
the sulphates, and carbon disulphide. 

Hydrogen Sulphide 

Hydrogen sulphide, H 2 S, is a gaseous compound of sulphur 
and hydrogen, and is sometimes called sulphuretted hydrogen 
or hydrosulphuric acid. It occurs in some volcanic gases and 
in the waters of sulphur springs. It is often found in the air, 
especially near sewers and cesspools, since it is one product 
of the decay of organic substances containing sulphur. It 
is one of the impurities of crude illuminating gas. 



338 INORGANIC CHEMISTRY 

Preparation. — The gas is prepared in the laboratory by 
the interaction of dilute acids and metallic sulphides; usually 
dilute hydrochloric acid and ferrous sulphide are used. 
When the acid is poured upon fragments of the sulphide, 
the gas is rapidly evolved without applying heat, and may 
be collected over water. The equation for the chemical 
change is — 

FeS + 2HC1 = H 2 S + FeCl 2 

Iron Hydrochloric Hydrogen Iron 

Sulphide Acid Sulphide Chloride 

Properties. — Hydrogen sulphide gas is colorless and has 
the odor of rotten eggs. It is poisonous. A little, even if 
diluted with air, often produces headache and nausea, and 
a large quantity of the gas is fatal. Care should be used 
in working with hydrogen sulphide, especially if the genera- 
tor is large. The dry gas is slightly heavier than air; a 
liter under standard conditions weighs 1.537 gm. It has 
been liquefied and solidified by the usual methods. Hydro- 
gen sulphide is soluble in water, one volume of water dis- 
solving about three volumes of the gas at the ordinary 
temperature; the dissolved gas can be completely removed 
by boiling the solution. The solution is called hydrogen 
sulphide water, and is often used instead of the gas in 
chemical experiments. The solution decomposes slowly, 
sulphur being deposited. 

Hydrogen sulphide gas is inflammable and burns with a 
bluish flame, forming water and sulphur dioxide, thus: — 

2 H 2 S + 3 2 = 2 S0 2 + 2 H 2 

Hydrogen Oxygen Sulphur Water 
Sulphide Dioxide 

If the supply of air is insufficient, combustion is incom* 
plete, water and sulphur being formed, thus : — 

2H 2 S + 2 = 2H 2 + 2S 

Hydrogen Oxygen Water Sulphur 

Sulphide 



SULPHUR AND ITS COMPOUNDS 339 

It is a reducing agent, and is used as such in chemical ex- 
periments. Even sulphuric acid is reduced by it, thus : — 



H 2 S0 4 


+ H 2 S = S0 2 + S + 


2H 2 


Sulphuric 
Acid 


Hydrogen Sulphur Sulphur 
Sulphide Dioxide 


Water 



Hydrogen sulphide gas dissolved in water gives a solution 
which has a feeble acid reaction, is neutralized by bases, 
and forms salts called sulphides. It dissociates only slightly, 
the ions being chiefly H + and HS" ; some S" ~ ions are formed. 

Composition of Hydrogen Sulphide Gas. — When metals are heated 
in dry hydrogen sulphide, metallic sulphides are formed; and the 
volume of hydrogen liberated by their decomposition is the same as 
the original volume of gas used. Since the hydrogen molecule con- 
tains two atoms (H 2 ), there must be two atoms of hydrogen in the 
hydrogen sulphide molecule. The vapor density of hydrogen sul- 
phide gas as found by experiment requires the molecular weight 
34 (approximately). Subtracting 2 for 2 H, the remainder (32) agrees 
well with the atomic weight of sulphur. Hence, there can be only 
one atom of sulphur in hydrogen sulphide, and the formula must 
be H 2 S. 

Sulphides may be regarded as salts of the weak acid, 
hydrosulphuric acid, though they are not always prepared 
directly from hydrogen sulphide by the substitution of a 
metal for its hydrogen. They may be produced by the 
direct union of sulphur and metals' (as in the case of iron 
and copper sulphides previously mentioned), by exposing 
the metal to the moist gas, or by the reduction of a sul- 
phate with carbon. A more common way is to precipitate 
them by passing the gas into solutions of metallic compounds, 
or sometimes, by adding hydrogen sulphide water to such 
solutions. Copper, tin, lead, and silver are rapidly tar- 
nished by the gas. Silverware, on this account, turns brown 
or black, especially in houses heated by coal or lighted by 
coal gas. The brown silver sulphide also coats silver spoons 



340 INORGANIC CHEMISTRY 

which are put into mustard or eggs. Lead compounds are 
blackened by this gas, owing to the formation of lead sul- 
phide, thus : — 

PbO + H 2 S ■ = PbS + H 2 

Lead Hydrogen Lead Water 

Oxide Sulphide Sulphide 

For this reason houses painted with " white lead" paint 
often become dark, and, similarly, oil paintings are dis- 
colored. The darkening (to brown or black) of the solu- 
tion of a lead compound is the customary test for hydrogen 
sulphide. 

Many sulphides have a brilliant color. Arsenious sulphide is 
pale yellow, cadmium sulphide is golden yellow, manganese sulphide 
is flesh colored, zinc sulphide is white, antimony sulphide is orange 
red. They vary in solubility. Most sulphides are insoluble in water. 
The sulphides of lead, silver, copper, and some other metals are in- 
soluble in dilute hydrochloric acid. The sulphides of iron, zinc, and 
some other metals are decomposed by dilute hydrochloric acid, but 
are precipitated if ammonium hydroxide is present. Sulphides of 
certain metals dissolve in water. Hence by precipitating metals 
under different conditions, groups of metals may be separated and 
subjected to further tests. The color often affords a ready means of 
detecting each sulphide. Hydrogen sulphide is thus a serviceable 
reagent in Qualitative Analysis. 

Oxides of Sulphur 

Sulphur Dioxide, S0 2 , is the common compound of sulphur 
and oxygen. It occurs in the gases of volcanoes, and to a 
slight extent in the atmosphere, since it is the usual product 
of the combustion of sulphur and sulphur compounds. 

Preparation. — When sulphur burns in air (or oxygen), 
sulphur dioxide is formed, thus : — ■ 

S + 2 = S0 2 

Sulphur Oxygen Sulphur 

Dioxide 



SULPHUR AND ITS COMPOUNDS 341 

It is also formed by roasting iron disulphide (iron pyrites) 
in the air, thus : — 

4 FeS 2 + 11 2 = 8 S0 2 + 2 Fe 2 3 

Iron Oxygen Sulphur Iron 

Disulphide Dioxide Oxide 

The foregoing reaction is utilized on a large scale in the 
commercial manufacture of sulphuric acid. Sulphur and 
carbon reduce sulphuric acid to sulphur dioxide, thus : — 

S + 2H 2 S0 4 = 3S0 2 + 2H 2 

Sulphur Sulphuric Sulphur 

Acid Dioxide 

C + 2 H 2 S0 4 = 2 S0 2 + C0 2 + 2 H 2 

Carbon Carbon 

Dioxide 

Two methods of preparation are used in the laboratory. 

(1) If copper and concentrated sulphuric acid are heated, 
a series of complex changes results finally in the evolution 
of sulphur dioxide. The equation is usually written : — 

Cu + 2H 2 S0 4 = S0 2 + CuS0 4 + 2 H 2 

Copper Sulphuric Sulphur Copper 
Acid Dioxide Sulphate 

(2) Sulphuric (or hydrochloric) acid dropped upon a sul- 
phite yields sulphur dioxide, thus : — 

Na 2 S0 3 + H 2 S0 4 = S0 2 + Na 2 S0 4 + H 2 

Sodium Sulphuric Sulphur Sodium 

Sulphite Acid Dioxide Sulphate 

The latter method is convenient for liberating a steady 
current of the gas. (See page 344.) 

Properties. — Sulphur dioxide gas has no color. It has 
a suffocating odor, being the well-known odor readily noticed 
when sulphur is burned. It will not burn in the air, nor 
will it support ordinary combustion. A burning taper or 
stick of wood is instantly extinguished by it, but finely 



342 INORGANIC CHEMISTRY 

divided metals, iron for example, burn in it. It is a heavy 
gas, the high density (2.2) allowing it to be readily collected 
by displacement of air, as in the case of chlorine. A liter 
of sulphur dioxide under standard conditions weighs 2.927 
gm. Its critical temperature is about + 155.5° C, and at 
a moderately low temperature it changes into a transparent, 
colorless liquid, which boils at — 8° C; if sufficiently cooled, 
it freezes into a transparent, icelike solid, which melts at 
about — 76° C. Liquid sulphur dioxide is a common com- 
mercial article. Sulphur dioxide gas is very soluble in 
water. At the ordinary temperature one volume of water 
dissolves about forty volumes of gas, but the solution loses 
it all by boiling. This solution is sour and reddens blue 
litmus, and contains sulphurous acid (see below). Moist 
sulphur dioxide bleaches vegetable coloring matters. A red 
or purple flower loses color in it. Silk, hair, straw, wool, 
and other delicate substances, which would be injured by 
chlorine, are whitened by sulphur dioxide. In some cases 
the color returns when the bleached article is exposed to 
the air for some time, and usually such bleached objects 
become yellow with age. The coloring matter is not wholly 
destroyed, but probably unites with the sulphur dioxide to 
form a colorless compound, which slowly decomposes. 

Uses. — Immense quantities of sulphur dioxide are used 
in the manufacture of sulphuric acid. The gas is also used 
to preserve meat and wines, to fumigate clothing and houses, 
in paper making, in tanning, in refining sugar, and in mak- 
ing acid sodium sulphite. Liquid sulphur dioxide is used 
in extracting glue and gelatine, and in various metallurgical 
processes. It absorbs considerable heat during evapora- 
tion, and was formerly used in some ice machines. 

The Composition of Sulphur Dioxide is based on the following: 
The gas formed by burning sulphur in a measured volume of oxygen 



SULPHUR AND ITS COMPOUNDS 343 

has the same volume as the oxygen itself. Hence there are as many 
molecules of .sulphur dioxide as there were of oxygen; that is, one 
molecule of sulphur dioxide contains two atoms of oxygen. Two 
atoms of oxygen weigh 32. But the molecular weight of sulphur 
dioxide calculated from an experimental determination of its vapor 
density is about 64. Subtracting 32 (i.e. 2x 16) for the oxygen (2 O) 
from this, there remains about 32 for sulphur. The atomic weight 
of sulphur is 32.07; hence sulphur dioxide contains only one atom 
of sulphur, and its composition is expressed by the formula SO2. 

Sulphur Trioxide, S0 3 , is formed by the direct union of 
sulphur dioxide and oxygen, though a little is produced 
when sulphur burns in air or in oxygen. The action is slow, 
but can be hastened by passing a mixture of sulphur dioxide 
and oxygen (or air) over hot platinum, or better over asbestos 
coated with platinum. Other substances also hasten the 
chemical change. At the ordinary temperature sulphur 
trioxide is a liquid, which boils at 46° C. and solidifies at 
15° C. to a white c^stalline mass. Another solid form 
(sulphur hexoxide, S 2 6 ), resembling asbestos, is known. 
When exposed to mois.t air, it fumes strongly, forming sul- 
phuric acid ; and when dropped into water, it dissolves with 
a hissing sound and evolution of heat. The equation for this 
chemical change is — 

S0 3 + H 2 = H 2 S0 4 

Sulphur Water Sulphuric 
Trioxide Acid 

Oxygen Acids and Salts of Sulphur 

Sulphurous Acid and Sulphites. — Sulphurous acid, H 2 S0 3 , 
is formed when sulphur dioxide dissolves in water. Sulphur 
dioxide is, therefore, sulphurous anhydride. The simplest 
equation expressing this fact is — 

S0 2 + H 2 = H 2 S0 3 

Sulphur Water Sulphurous 
Dioxide Acid 



344 INORGANIC CHEMISTRY 

This acid is known only in solution, and resembles car* 
bonic acid in this respect. It is unstable and decomposes 
readily into sulphur dioxide and water. It gradually forms 
sulphuric acid by combining with oxygen from the air and 
very rapidly by interaction with oxidizing agents, such 
as potassium permanganate. Sulphurous acid is dibasic, 
and forms two classes of salts, the normal and acid sul- 
phites. They are reducing agents, and yield sulphur dioxide 
when treated with acids. It will be recalled in this con- 
nection that a convenient method of preparing sulphur 
dioxide is the decomposition of normal sodium sulphite 
(Na 2 S0 3 ) by sulphuric or hydrochloric acid. The sulphite 
first forms sulphurous acid (H 2 S0 3 ), and the unstable acid 
decomposes into water and sulphur dioxide. The following 
equations express these two reactions : — 

Na 2 S0 3 + H 2 S0 4 = H 2 S0 3 + Na 2 S0 4 

Sodium Sulphuric Sulphurous Sodium 

Sulphite Acid Acid Sulphate 

H 2 so s = so 2 + H 2 

Sulphurous Sulphur Water 

Acid Dioxide 

Acid sodiuni sulphite (HNaS0 3 ), often called bisulphite of 
soda, is the anticolor used to remove the excess of chlorine 
from bleached cotton cloth. It is also used in brewing, 
tanning, and in the manufacture of starch, sugar, and 
paper. Acid calcium sulphite (CaH 2 (S0 3 ) 2 ), also called 
bisulphite of calcium, is prepared by passing sulphur di- 
oxide into milk of lime, and is used in paper making. 

Sulphurous acid is rather weak. In terms of the theory of 
electrolytic dissociation it dissociates only to a compara- 
tively slight extent, the ions being chiefly H + and HS0 3 "; 
the latter ion, however, dissociates to some degree into H + 
and S0 3 ~~. Solutions of the acid sulphites of sodium and 
potassium have an acid reaction, owing to the slight disso- 



SULPHUR AND ITS COMPOUNDS 345 

ciation of the HS0 3 -ion into the ions H + and S0 3 ~~. Solu- 
tions of the normal sulphites of sodium and potassium are 
alkaline. This fact is briefly explained in terms of the 
theory of electrolytic dissociation as follows: The S0 3 -ions 
from the sulphite tend to form the more stable HS0 3 -ions 
by combining with the H-ions furnished by the appreciably 
dissociated water; hence OH-ions are left in the solution 
and give it an alkaline reaction. As already stated, cases 
of double decomposition like the one just cited in which 
water is one factor are illustrations of hydrolysis (see 
Hydrolysis, Chapter X). 

Sulphuric Acid, H 2 S0 4 , is found in the waters of a few 
rivers and mineral springs. It is manufactured in enor- 
mous quantities and used for many purposes. 

Sulphuric acid was doubtless known to the Arabian alchemists 
living in the tenth century. It was definitely mentioned by Basil 
Valentine in the fifteenth century, who described its preparation 
by heating a mixture of iron sulphate (green vitriol) and sand. The 
product, an oily liquid, was called oil of vitriol, a name now often 
applied to commercial sulphuric acid. 

Sulphuric acid is manufactured by two processes, the lead 
chamber process and the contact process. In both processes 
sulphur dioxide is oxidized to sulphur trioxide, which is 
transformed into sulphuric acid by combination with water. 
A skeleton equation, so to speak, may be written thus : 



2 SO* 


+ o 2 


+ 


2H 2 


= 


2 H 2 S0 4 


Sulphur 


Oxygen 




Water 




Sulphuric 


Dioxide 










Acid 



The oxidation is accomplished in two ways. In the older 
or lead chamber process oxides of nitrogen are used. In 
the newer or contact process a suitable catalytic agent, usually 
finely divided platinum, is employed. 



346 INORGANIC CHEMISTRY 

In the chamber process sulphur dioxide, air, steam, and 
oxides of nitrogen are introduced into large lead chambers. 
These gases interact and produce sulphuric acid, which col- 
lects on the floors of the lead chambers. The chemical 
changes involved in this process are complex and variable. 
The main reactions are : — 

2 S0 2 + N2O3 + 2 + H 2 = 2 S0 2 (OH)(ONO) 

Nitrosyl-sulphuric 
Acid 

2 S0 2 (OH)(ONO) + H 2 = 2 H 2 S0 4 + N 2 3 

The nitrogen trioxide (N 2 3 ) is apparently an' essential 
factor, though according to some authorities the change 
may be due to a mixture of nitric oxide (NO) and nitrogen 
peroxide (N0 2 ). Theoretically a small quantity of oxides 
of nitrogen is needed to form an unlimited amount of sul- 
phuric acid. Some is lost, however, and must be replaced. 
This is done by putting nitric acid into the top of the Glover 
tower or by injecting it into the chambers. 

The details of the lead chamber apparatus are shown in Figure 
53. There are three main parts : (a) the furnace for producing 
sulphur dioxide, (b) the lead chambers together with the Glover 
and Gay-Lussac towers for changing the gaseous mixture into sul- 
phuric acid, and (c) the concentrating apparatus. The manufac- 
ture is conducted as follows: (1) Sulphur or iron disulphide (iron 
pyrites, FeS 2 ) is burned in a furnace with enough air to change the 
sulphur into sulphur dioxide, and to furnish the proper amount of 
oxygen for later changes. In some works the furnace is provided 
with " niter pots " to produce nitric acid vapor. (2) The mixture 
of sulphur dioxide, oxides of nitrogen, and air passes from the fur- 
nace into the bottom of the Glover tower. This is a tall tower 
filled with small stones over which flow two streams of sulphuric 
acid, one dilute and the other containing nitrogen dioxide (from 
the Gay-Lussac tower acid). These acids cool the ascending gases ; 
the dilute acid is deprived of water and the tower acid of nitro- 
gen dioxide. Hence, concentrated acid flows out of the bottom of 



SULPHUR AND ITS COMPOUNDS 



347 



the Glover tower, 
while from the 
top sulphur diox- 
ide, oxides of ni- 
trogen, steam, and 
air pass into the 
first lead chamber. 
Here nitric acid 
is introduced, as 
well as steam. 
The main chemi- 
cal changes occur 
in this and in the 
second chamber. 
These chambers 
are huge boxes 
of sheet lead 
supported on a 
wooden frame- 
work, The un- 
used gases pass on 
into the bottom 
of the Gay-Lussac 
tower. This tower 
is filled with coke 
or earthenware 
over which flows 
concentrated sul- 
phuric acid (ob- 
tained from the 
Glover tower), 
which absorbs the 
unused nitrogen 
dioxide. The ox- 
ide, as stated 
above, is liber- 
ated again in the 
Glover tower. 
(3) The acid pro- 
duced in the cham- 




Fig. 53. — Apparatus for chamber process. 



348 



INORGANIC CHEMISTRY 




bers contains from 60 to 70 per cent 
of H2SO4. It is concentrated into 
commercial sulphuric acid, which 
contains about 96 to 98 per cent, by 
evaporation, first in lead-lined pans 
and finally in a platinum vessel. 

In the contact process sul- 
phur dioxide and air, carefully 
purified and heated to about 
400° C, are brought in contact 
with a catalytic agent (see page 
185), which is usually finely 
divided platinum. The sulphur 
dioxide is oxidized to sulphur 
trioxide, thus : — 

2 S0 2 + 2 = 2 S0 3 

The sulphur trioxide is con- 
ducted into sulphuric acid con- 
taining a little water with which 
the trioxide combines, thereby 
forming sulphuric acid, thus : — • 

S0 3 + H 2 = H 2 S0 4 

The contact process apparatus 
is shown in Figure 54. The blower 
A forces air into the burner B, where 
the suphur dioxide is formed by 
burning iron pyrites (FeS 2 ) or sul- 
phur. The gases pass into C, where 
they are freed from sulphur dust 
and other solid impurities. The 
gases, cooled by the pipe D, are 
further cleaned in the scrubbers, 
which contain coke wet with water 
(E) and with sulphuric acid (F), 



SULPHUR AND ITS COMPOUNDS 349 

and then freed from arsenic compounds in G. Traces of such com- 
pounds " poison " the platinum and stop the formation of sulphur 
trioxide. The purified gases (mainly sulphur dioxide) then enter 
the mixer and heater H . Here a large excess of air is introduced 
and the whole mixture is heated to about 400° C. The mixture of 
sulphur dioxide and air next passes into the contact chamber /. 
Here the gases come in contact with the catalytic agent and form 
sulphur trioxide. The catalytic agent, if platinum, consists of 
asbestos fibers coated with a very thin layer of metallic platinum 
and is spread out on plates or mixed with porous material in order 
to provide a large contact surface. The sulphur trioxide finally 
passes into an absorber J partly filled with sulphuric acid contain- 
ing 2 to 3 per cent of water. In this liquid, the combination with 
water takes place readily, water being added to maintain the required 
concentration in the absorber. 

Properties of Sulphuric Acid. — Sulphuric acid is an oily 
liquid, colorless when pure, though often brown from the 
presence of charred organic matter, such as dust and straw. 
The specific gravity of the commercial acid is about 1.83. 
When sulphuric acid is mixed with water, much heat is 
evolved. The acid should always be poured into the water, 
otherwise the intense heat may crack the vessel or spatter 
the hot acid. Sulphuric acid combines with water and forms 
hydrates ; a rather stable one has the formula : — 

H 2 S0 4 . H 2 0. 

The tendency to absorb water is shown in many ways. 
The concentrated acid absorbs moisture from the air and 
from gases passed through it. It is often used in the labora- 
tory to dry gases, since it is not volatile at the ordinary 
temperature. Wood, paper, sugar, starch, cotton cloth, 
and many organic substances are blackened by sulphuric 
acid. Such compounds contain hydrogen and oxygen in 
the proportion to form water; these two elements are ab- 
stracted and carbon alone remains. Similarly, sulphuric 



350 INORGANIC CHEMISTRY 

acid withdraws water from the flesh, making painful 
wounds. 

Sulphuric acid boils at about 338° C, but begins to give 
off fumes of sulphur tri oxide at about 150° C. It oxidizes 
carbon and sulphur to carbon dioxide and sulphur dioxide, 
owing to its instability when heated with these elements. 
Metals decompose it, yielding various products, such as 
hydrogen, sulphur dioxide, or hydrogen sulphide. It com- 
bines directly with ammonia gas, thus : — 



2NH 3 


+ 


H 2 S0 4 = 


= (NH^SO, 


Ammonia 




Sulphuric 


Ammonium 


Gas 




Acid 


Sulphate 



Its interaction with salts, such as chlorides, nitrates, and 
sulphites, which results in the liberation of the correspond- 
ing acid, has already been discussed. A solution of sul- 
phuric acid in water contains hydrogen ions (H + ), monova- 
lent ions (HSO4-), and divalent ions (S0 4 ""~), depending upon 
the concentration. Dilute solutions contain an abundance 
of S0 4 -ions (S0 4 — ) and H-ions (H + ). 

Uses of Sulphuric Acid. — Sulphuric acid is one of the 
most important substances. Directly or indirectly it is 
used in hundreds of industries upon which the comfort, 
prosperity, and progress of mankind depend. It is used 
in the manufacture of other mineral acids and many organic 
acids. It is essential in one process of manufacturing 
sodium carbonate. Enormous quantities are consumed in 
making artificial fertilizers, alum, nitroglycerin, glucose, 
phosphorus, dyestuffs, and in various parts of such funda- 
mental industries as dyeing, bleaching, electroplating, refin- 
ing, and metallurgy. 

Sulphates. — Sulphuric acid is dibasic and forms two 
classes of salts, — the normal sulphates, such as Na 2 SC>4, 



SULPHUR AND ITS COMPOUNDS 351 

f 

and the acid sulphates, such as HNaS0 4 . The normal sul* 
phates are rather stable salts ; many yield sulphur trioxide 
when heated to a high temperature. The acid salts lose 
water when heated, yielding in addition salts called pyro- 
sulphates (see below under pyrosulphuric acid). Most sul- 
phates are soluble in water, only the sulphates of barium, 
strontium, and lead being insoluble, while calcium sulphate 
is slightly soluble. Important sulphates are calcium sul- 
phate (gypsum, CaS0 4 .2H 2 0), barium sulphate (heavy 
spar, BaS0 4 ), zinc sulphate (white vitriol, ZnS0 4 ), copper 
sulphate (blue vitriol or bluest-one, CuS0 4 ), iron sulphate 
(green vitriol, copperas, ferrous sulphate, FeSO^, sodium 
sulphate (Glauber's salt, Na 2 S0 4 ), and magnesium sulphate 
(Epsom salts, MgS0 4 , and kieserite, MgS0 4 .H 2 0). Sul- 
phates are widely used as medicine and in many industries. 
They are described more fully under the individual metals. 

The test for sulphuric acid or a soluble sulphate is the 
formation of white, insoluble barium sulphate upon the 
addition of barium chloride solution. An insoluble sulphate, 
such as calcium sulphate, when fused on charcoal is reduced 
to a sulphide, w^hich blackens a moist silver coin. The 
blackening is due to the formation of silver sulphide. 

The usual test, as already stated, is applicable to both the 
acid and soluble sulphates because their solutions contain 
S0 4 -ions. When a solution containing barium ions is added, 
the ions Ba + + and S0 4 " unite and form barium sulphate, 
which is insoluble in water and is not (like some other 
barium salts) decomposed by the common acids. The ionic 
equation for the reaction is : — 

Ba + + + SO" — ^BaS0 4 

Fuming Sulphuric Acid, HgSoO;, is made by adding sul- 
phur trioxide to sulphuric acid, or by heating moist ferrous 
sulphate. It is sometimes called Nordhausen sulphuric 



352 INORGANIC CHEMISTRY 

acid. It is a thick, brown liquid, which fumes strongly in 
the air, owing to the escape of oxides of sulphur. It is 
used in gas analysis to absorb ethylene and other illuminants, 
and in dyeing to dissolve indigo. If the fuming acid is 
cooled to 0° C, crystals separate; they are called pyrosul- 
phuric acid or disulphuric acid. Its salts are the pyro- 
sulphates. 

Other Sulphur Compounds 

Sodium Thiosulphate, Na 2 S 2 3 , is a salt of an unstable 
acid. It is sometimes incorrectly called sodium hyposul- 
phite, or simply "hypo." It is a white, crystalline solid, 
very soluble in water. The solution, used in excess, dis- 
solves the halogen compounds of silver, i.e. AgCl, AgBr, 
Agl; hence its extensive use in photography. (See Photog- 
raphy.) It is used in dyeing, and in chemical analysis 
(for determining the amount of free iodine in a solution). 

Carbon Disulphide (or Carbon Bisulphide, CS 2 ) when pure 
is a clear, colorless liquid, with an agreeable odor; the com- 
mercial substance is yellow and has an offensive odor. It is 
poisonous. Carbon disulphide is an endothermic com- 
pound, i.e. its synthesis from its elements is accompanied by 
absorption of heat. The thermal equation for the forma- 
tion of carbon disulphide is — 

C + S 2 = CS 2 - 28,700 cal. 

Carbon Sulphur Carbon 

Disulphide 

Like all endothermic compounds it is relatively unstable 
and can be exploded by mercuric fulminate. Ordinarily it 
can be handled without danger of explosion. It is volatile 
and extremely inflammable, the equation for its combustion 
being — 

CS 2 + 3 2 = C0 2 + 2 S0 2 

Carbon Oxygen Carbon Sulphur 

Disulphide Dioxide Dioxide 



SULPHUR AND ITS COMPOUNDS 



353 



This liquid is practically insoluble in water. It dissolves 
rubber, gums, fats, resins, iodine, camphor, and some forms 
of sulphur. It is a highly refracting liquid, and hollow 
glass prisms filled with it are used to decompose light. As 
a solvent it is used to dissolve pure rubber in the manu- 
facture of rubber cement. It is also used to kill insects on 
both living and dried plants (e.g. in museums), and to ex- 
terminate burrowing animals, such as moles and wood- 
chucks. Many oils, waxes, and 
greases are extracted by carbon 
disulphicie. 

Until recently carbon clisulphide 
was manufactured by passing sul- 
phur vapor over red-hot coke or 
charcoal in iron or earthenware 
retorts. The product required 
laborious purification. It is now 
manufactured by an electrothermal 
process in a furnace somewhat like 
that shown in Fig. 55. Several 
groups of carbon electrodes (E, E) 
are set into the base of a furnace, 
coke is packed loosely between them; 

Sulphur is put into Z and partly Fig. 55. — Furnace for the 

surrounds the electrodes. The body manufacture of carbon di- 
of the furnace is filled with char- sup 
coal (C). Sulphur is introduced at suitable points (S, S, S), 
and coke is fed in through K, K. When the current passes, 
the sulphur melts and unites with the heated carbon above 
the electrodes. The vapors of the resulting carbon disulphide 
escape from the top of the furnace through P, and are 
condensed in a special apparatus. 




354 INORGANIC CHEMISTRY 

Problems and Exercises 

1. Sulphur is burned in 10 1. of air containing 21 per cent of 
oxygen by volume and measured at 22° C. and 767 mm. Calculate 

(a) the weight of sulphur burned, (6) the volume of oxygen con- 
sumed, (c) the weight of sulphur dioxide produced. 

2. Suppose a manufacturer of H 2 S0 4 starts with 100 tons of 
sulphur and obtains the theoretical yield in each case: (a) What 
weight of oxygen is needed to burn the sulphur to sulphur dioxide ? 

(b) What additional weight of oxygen to convert the sulphur dioxide 
to sulphur trioxide? (c) What weight of water to convert the sul- 
phur trioxide to sulphuric acid ? (d) And how much sulphuric acid 
is obtained? 

3. What weight of pure H 2 S0 4 can be manufactured from 150 
tons (2000 lb. each) of iron pyrite containing 92 per cent of FeS 2 ? 

4. What weight of sulphuric acid (sp. gr. 1.8354) is contained 
in a cylindrical tank 20 m. long and 1.3 m. in diameter? 

5. Calculate the atomic weight of sulphur from the following 
data: (a) In BaS0 4 , Ba = 58.85, S = 13.73, = 27.42; (b) 10 gm. of 
silver yield 11.4815 gm. of silver sulphide; (c) 2 gm. of lead yield 
2.9284 gm. of lead sulphate. (Use exact atomic weights.) 

6. What weight of carbon disulphide is formed by the inter- 
action of (a) carbon and 200 kg. of sulphur and (b) sulphur and 200 
kg. of carbon? 

7. Express the following reactions by volumetric equations: 
(a) Sulphur dioxide and oxygen form sulphur trioxide, (b) hydrogen 
sulphide and oxygen form sulphur dioxide and water (vapor), 

(c) carbon disulphide (vapor) and oxygen form sulphur dioxide and 
carbon dioxide. 

8. Write the formulas of (a) the acid sulphates, (b) the sulphites, 
(c) the acid sulphites, and (d) the sulphates of Ba, Ca, K, Na. 

9. Write the formulas of the sulphides of NH 4 , Al, Ba, Ca, Cr, 
Cu(ic), Hg(ic), Ag, Sn(ous), Zn. 

10. Complete and balance the following equations: (a) H 2 S0 4 

+ H 2 S = S0 2 + + H 2 ; (b) SrC0 3 + H 2 S0 4 = SrS0 4 *+ 

+ H 2 ; (c) CdCl 2 + H 2 S = HC1 + - — ; (d) Pb(N0 3 ) 2 

+ = PbS0 4 + HN0 3 ; (e) NH 3 + = (NH 4 ) 2 S0 4 . 

11. If a sample of carbon disulphide has a specific gravity of 
1.25, what volume can be made from 2 metric tons of pure sulphur? 

12. Calculate the solubility product of lead sulphate if the ioni- 
zation is 100 per cent and the molar solubility is .00013. 



CHAPTER XIX 
Classification of the Elements 

Introduction. — In the preceding chapters emphasis has 
been laid on the individual elements, especially oxygen, 
hydrogen, nitrogen, carbon, chlorine, and sulphur. These 
elements differ from each other in many ways, and if this 
diversity prevailed among all the eighty elements, it would 
be difficult to proceed very far with the study of chemistry. 
Fortunately the elements are not independent. Certain 
ones are so similar in their chemical relations that they can 
be put into the same class. Several characteristic classes 
have been formed. The arrangement into groups or classes 
not only simplifies the study of the elements, but reveals 
many fundamental relations. 

Classification of the Elements. — Many attempts have been 
made to classify the elements. About the time of Lavoisier 
(1743-1794) they were roughly divided into metals and non- 
metals. Those elements were called metals which were 
hard, lustrous, heavy, and good conductors of heat, while 
the others were called non-metals. This classification 
proved to be misleading as additional elements w r ere dis- 
covered whose properties did not harmonize with the prin- 
ciple of division. It is used at the present time, however, 
because many common elements fall readily into one of 
these classes, as shown in the lists of metals and non-metals 
given in Chapter X. 

Classification according to acid and basic properties pre- 
vailed for a time. But it was abandoned largely because 
such a basis of division excluded elements exhibiting both 

355 



356 INORGANIC CHEMISTRY 

acid and basic properties, such as zinc and chromium. The 
elements have also been classified according to their valence 
into six or seven groups (the mono-, di-, tri-, etc.). But 
this plan was given up partly because of so many cases of 
variable valence, but mainly on account of the indiscrim- 
inate character of the groups. For example, elements so 
unlike as sodium and chlorine were in the same group. 

About 1828 striking resemblances between certain ele- 
ments were pointed out, and several groups or families 
were suggested. For example : — 

Lithium Selenium Calcium , Nitrogen 

Sodium Sulphur Strontium Phosphorus 

Potassium Oxygen Barium Arsenic 

This classification was arbitrarily based on selected physical 
and chemical properties. It was interesting but incom- 
plete, because it emphasized resemblances and overlooked 
differences; that is, the basis of comparison was not broad 
enough. 

The first actual progress began to be made about 1850, 
when chemists became deeply interested in the significance 
of atomic weights. Dumas and others pointed out certain 
striking numerical relations between the atomic weights of 
related elements. Thus> the atomic weight of sodium is 
half the sum of the atomic weights of lithium and potassium. 

Li = 7, Na = 23, K = 39. 1±^ = 23 

2 

The same is true of phosphorus, arsenic, and antimony. 

P = 31, As = 75, Sb - 120. 31 + 120 = 75.5 

2 

A number of such families was found, but this method of 
classification was not comprehensive, nor did it entirely 
eliminate chance selections. 

The existence of other relations similar to those just 




MENDELEJEFF 

1834-1907 



CLASSIFICATION OF THE ELEMENTS 357 

cited, a deep desire to obtain more accurate atomic weights, 
and a growing interest in the properties of the elements 
themselves focused the attention of chemists about 1860 
upon the relation of properties to atomic weights. Several 
conditions fostered this principle of classification. One was 
the atomic weight determinations of Stas, whose masterly 
work yielded very exact atomic weights. Another was the 
acceptance by most chemists of a uniform table of atomic 
weights. A third was the vast accumulation of facts about 
the elements and their compounds. Chemists were ready 
for a new and broader classification of the elements. 

The Periodic Classification of the Elements. — Previous 
to 1869 no classification included all the elements. In that 
year the Russian chemist Mendelejeff published an arrange- 
ment known as the periodic classification of the elements. 
This classification revealed a new relation between the 
properties of the elements and their atomic weights, and is 
very helpful in studying the chemical elements. Men- 
cielejefTs scheme, slightly modified to conform to subsequent 
discoveries, is substantially as follows : — 

If all the elements are arranged in the order of their in- 
creasing atomic weights, a series results in which similar 
or closely related elements occur at regular intervals. That 
is, the series breaks up into several periods, and hence the 
system of classification is called periodic. If the series is 
divided into these periods and the periods placed below 
each other, a table is secured in which the perpendicular 
columns consist of groups of similar elements, i.e. elements 
which have similar relations and form analogous compounds 
which have similar properties. Such a table is shown on 
the following page. In each group the elements have been 
subdivided into families in order to emphasize the mutual 
relationship among closely allied elements. 



358 



INORGANIC CHEMISTRY 



H 

to 

W 

H 

h 

O 

to 

o 

H 






Q 
O 

s 

Ph 



Ph 



o 
PS 








t* OO b- 
00 tO 3S 

O OO GO 

<» -~ C 

^ ^5 O 




h- Cl t- 
-h CM «c 

© © o 

1* 1* If 

3 ,C r- 




© ^ CM 
I'll 
I J J 

O A Ph 






Ph 

O 
Pi 
O 




© 


CO 
II 

5 


CO 

Tti 

l*"*j 


© 
© 

7 




© 
©' 

jl 








a, 
ta 

O 
Ph 

O 




o 

CO 

T 


o 

CO 
00 


© 
©J 

II 

Fh 

Q 


Oa 

0) 
GO 


© 
© 
© 

II 

o 


1© 
*-* 
CM 

r-t 

H 


© 

GO 




CM 

GO* 

II 


Pi 
O 




© 

7 


o 

tH 

CO 

Cm 


© 


© 
© 

II . 

CO 


CO 

© 
II 

o 


CM 

©' 
CM 

£i 
OD 


CO 

H 


© 
go" 
© 

CM 

1! 

s 




o 




© 
7 


CO 
GO 

35 


GO* 

H 


©S 


© 
© 
© 

S-c 

N 


GO 

a 
cc 


i ?2 

Q 3 


© 

CM 
CM 


CM* 

CO 

X! 






© 
CO 


T—t 

CM 

Jl 

< 


T-t 

II 

O 


©" 

© 

II 

O 


GO 
II 
K- 1 


CO 

7 


o 

© 

CO 

If 

c3 


© 

3 

CM 

II 

H 




u 

o 




| 


CO 

tH 

CM 

II 

be 
3 


© 

II 

O 


CO 

© 


CO 
© 

GO 
jl 
CO 


© 

CM 
O 


CO 
CO 

cq 


© 
© 

CM 

II 

be 


© 
©' 
CM 
CM 

II 

C5 

M 


O 


GO 
O 

© 


© 
5© 

3 


o 
© 

CO 
CM 

II 


© 

© 

CO 


IO 

CO* 

5 


to 

GO 
II 

Ph 


co 

GO 
CC 

< 


CO 

cm" 

OO 

7 

o 


CM 

© 

■3 




o 

p- 

o 

tf 

o 




II 


©» 

CM 

|25 


CO 

GO 
CS 

CO 

II 

«1 




CM 

cm' 

CO 

J 

►2 




CM 

CO 

7 

X 




**? 
c4 

CM 

CM 


a 
o 

Ph 


H 


N 


CO 


«■ 


10 


© 


r- 


QO 


OS 


© 

H 



CLASSIFICATION OF THE ELEMENTS 359 

From the table it is seen that the elements fall naturally 
into two large general classes called groups and periods. 
Those elements in the same vertical column belong to the 
same natural group, while those in the same horizontal row 
belong to the same period. Selecting from each group the 
important elements, we have the following distinct families: 
Group 0. Inert elements or argon family — helium, 

neon, argon, krypton, xenon, niton. 
Group I. Alkali metals or sodium family — lithium, 
sodium, potassium. 
Univalent heavy metals or copper family — 
copper, silver, gold. 
Group II. Alkaline earth metals or calcium family — 
calcium, strontium, barium, radium. 
Bivalent heavy metals or zinc family — mag- 
nesium, zinc, cadmium, mercury. 
Group III. Boron family — boron. 

Earth metals or aluminium family — alu- 
minium. 
Group IV. Tetravalent non-metals or carbon family — 
carbon, silicon. 
Tetravalent metals or tin family — tin, lead. 
Group V. Pentavalent non-metals and metals or nitrogen 
family — nitrogen, phosphorus, arsenic, anti- 
mony, bismuth. 
Group VI. Hexavalent metals or chromium family — 
chromium, molybdenum, tungsten, uranium. 
Hexavalent non-metals or oxygen family — 
oxygen, sulphur, selenium, tellurium. 
Group VII. Manganese family — manganese. 

Halogen elements or chlorine family — fluo- 
rine, chlorine, bromine, iodine. 
Group VIII. Iron family — iron, cobalt, nickel. 
Platinum family — platinum. 



360 INORGANIC CHEMISTRY 

In several of these families the similarity is marked, 
especially between the alkali metals and also the halogens ; 
in some cases the resemblance is not very striking except 
in a limited number of properties. So far, our treatment 
has ignored these groups, brief references only having been 
made to the argon family. In the succeeding pages, how- 
ever, the different groups will be emphasized both as to 
their individual members and their collective properties. 

As stated above, the elements in the same horizontal row 
belong to the same period. The periodic variations of the 
properties of certain typical elements may be illustrated by 
the first (I) and second (II) periods. Ignoring the argon 
family which is somewhat anomalous, and beginning with 
lithium, the general chemical properties vary regularly with 
increasing atomic weight. The metallic character typified by 
lithium gradually diminishes through beryllium and boron; 
while the feeble non-metallic character typified by carbon 
increases through nitrogen and oxygen until fluorine is 
passed and sodium is reached ; here the metallic character 
reappears. Proceeding onward from sodium, the same 
decrease of basic and increase of acid properties is noticed 
until potassium is reached, and here again the marked 
metallic character reappears, There is no sudden change 
in properties until we pass from one period to the next. 
Thus, fluorine at the end of the second period forms a 
strong acid, but sodium at the beginning of the third period 
forms a strong base. Similarly, chlorine is strongly acidic, 
but potassium, which is the first metal in the next period, 
is markedly basic ; chlorine is a typical non-metal, while 
potassium is a typical metal. Not all the periods are as 
typical as those just cited, but in many cases the progres- 
sive change in chemical properties is too obvious to ascribe 
to mere chance. Many of these relations will be pointed out 
in the following chapters. Indeed it is hardly possible to 



CLASSIFICATION OF THE ELEMENTS 361 

appreciate the full significance of the periodic classification 
until most of the elements have been studied. 

Periodic Law. — The relation between properties and 
atomic weights which brings about this periodic variation is 
general, and is often summarized in the form of a law known 
as the periodic law, thus : — 

The 'properties of the elements are periodic functions of their 
atomic weights. 

Function here means the exhibition of some special rela- 
tion, viz. that of properties to atomic weight. Interpreted 
freely, the law means: (1) properties and atomic weight are 
related, they depend upon each other; and (2) this relation 
is exhibited repeatedly at regular intervals. 

Conclusion. — An examination of the table on page 358 
shows some imperfections in the periodic classification. 
For example, there are gaps. These probably correspond 
to elements not yet discovered or fully investigated. Three 
• such gaps, which were in the original table, have been filled. 
When Mendelejeff proposed his arrangement, he predicted 
the discovery of three elements having definite properties. 
These elements — gallium, scandium, and germanium — 
have since been discovered, and now occupy their predicted 
places in the table. Possibly other gaps will be filled by 
newly discovered elements. Several elements do not fall 
into their proper places. Thus, the atomic weight of argon 
indicates that this element should exchange places with 
potassium. Similarly the positions of iodine and tellurium 
should be reversed ; their properties, however, necessitate 
the present places. Hydrogen has no appropriate place in 
the series. 

The periodic classification, although imperfect in some 
particulars, simplifies the study of chemistry, and will be 
utilized in its larger aspects in many of the following chapters. 



CHAPTER XX 
Fluorine — Bromine — Iodine 

Fluorine, bromine, and iodine, together with chlorine, 
constitute a family in the seventh (VII) periodic group, 
often called the halogens. They have similar chemical prop- 
erties, and form analogous compounds which likewise have 
similar properties, differing mainly in degree. 

Halogen means " sea-salt producer." It is applied to this 
group of elements because they form salts which are found in 
sea water and resemble sodium chloride (common salt or sea 
salt). Chlorides, bromides, fluorides, and iodides are called 
halides. The Greek word for salt, hals, suggested these 
descriptive terms. 

Chlorine was fully discussed in Chapter XII. 

Fluorine 

Occurrence. — Fluorine is the most active of all the halogen 
elements, and is never found free in nature. It occurs 
abundantly in combination with calcium as fluor spar or 
calcium fluoride (CaF 2 ). Other native compounds are 
cryolite (Na 3 AlF 6 ) and apatite (Ca 5 F(P0 4 )3). Minute quan- 
tities of combined fluorine are found in bones and blood, 
in the enamel of the teeth, and in sea and some mineral 
waters. 

The Isolation of Fluorine was accomplished in 1886 by 
Moissan, though many unsuccessful attempts had previously 
been made. He decomposed hydrofluoric acid by electricity, 
and collected the liberated fluorine. The achievement was 

362 




MOISSAN 

1852-1907 



FLUORINE 



363 



attended with tremendous difficulties, owing to the intense 
activity of the fluorine and its corrosive properties. 

The essential parts of the apparatus used by Moissan are shown 
in Figure 56. The U-tube, made of an alloy of platinum and iridium, 
is provided with tightly fitting stoppers of 
fluor spar (S, S). Through the stoppers pass 
the electrodes (E, E), of platinum-iridium, 
held in place by screw caps (C, C). Side 
tubes (7\ T) allow the liberated gases (fluorine 
and hydrogen) to be drawn off separately 
through platinum delivery tubes. Perfectly 
dry hydrofluoric acid is put into the U-tube, 
and dry acid potassium fluoride (HKF2) is 
added to enable the solution to conduct the 
current — liquefied hydrofluoric acid itself 
being a non-conductor. The U-tube is then 
cooled to a low temperature (— 23 to — 50° C), 
and on passing a current through the solution 
fluorine is evolved at the positive electrode 
(anode) and hydrogen at the negative elec- 
trode (cathode). The fluorine freed from 
hydrofluoric acid vapor was collected by 

Moissan at first in a platinum tube with a thin fluor spar plate closing 
each end, so he could look inside and examine the gas. Later he 
found that the electrolysis could be performed in a copper U-tube 
and pure dry fluorine could be collected in a glass tube. 




Fig. 56. — Moissan *s ap- 
paratus for preparing 
fluorine. 



Properties. — Fluorine is a gas having a greenish yellow 
color, though lighter and more yellowish than chlorine. 
The critical temperature is — 120° C. Subjected to pressure 
and a sufficiently low temperature, it condenses to a very 
pale yellow liquid, which boils at about — 187° C. The 
vapor density shows that the molecular weight is about 38 ; 
hence, each molecule contains two atoms and fluorine gas 
has the formula F 2 (the atomic weight being 19). 

Chemically fluorine is intensely active. Hydrogen, bro- 
mine, iodine, sulphur, phosphorus, carbon, silicon, and boron 



364 INORGANIC CHEMISTRY 

take fire in it, but oxygen, nitrogen, and argon do not unite 
with it. Most metals burn in it, forming fluorides. Gold 
and platinum are not attacked by it below red heat. Copper 
becomes coated with copper fluoride, which protects the 
metal, so that copper vessels may be used as fluorine gen- 
erators and reservoirs. Moissan used a copper U-tube to 
prepare large volumes. Water is decomposed by it at 
ordinary temperatures, while hydrocarbons are instantly 
decomposed, hydrofluoric acid and carbon fluorides being 
the important products. 

Hydrofluoric Acid is the compound of fluorine correspond- 
ing to hydrochloric acid. It is prepared by the interaction of 
a fluoride and concentrated sulphuric acid, just as hydro- 
chloric acid is prepared from a chloride and sulphuric acid. 
Calcium fluoride is usually used, and the experiment is per- 
formed preferably in a lead dish. The chemical change is 
represented thus : — 



CaF 2 


+ 


H 2 S0 4 = 


2HF + 


CaS0 4 


Calcium 
Fluoride 




Sulphuric 
Acid 


Hydrofluoric 
Acid 


Calcium 
Sulphate 



Pure hydrofluoric acid is a colorless liquid which boils at 
about 19° C. It is very volatile and is readily transformed 
into a colorless gas (hydrogen fluoride), which fumes in the 
air and dissolves in water, the aqueous solution being the 
commercial hydrofluoric acid. Both gas and solution are 
dangerous substances. The gas is extremely poisonous, and 
the liquid, if dropped on the skin, produces terrible sores. 
Owing to its corrosive action hydrofluoric acid is preserved 
and sold in hard rubber, lead, or wax bottles. The vapor 
density of hydrofluoric acid gas varies with the temperature. 
From 19° to 30° C. the vapor density indicates a molecular 
weight of 40, while at 88° C. the molecular weight is 20. 
The atomic weight of fluorine is 19. Hence at the lower 



BROMINE 365 

temperature the formula of the gas is H 2 F 2 and at the higher 
temperature it is HF. 

Hydrofluoric acid interacts with metals, thereby liberating 
hydrogen and forming fluorides. It also interacts with bases 
and metallic oxides and forms fluorides. There are two 
classes of fluorides — normal and acid. Thus, calcium 
fluoride (CaF 2 ) is a normal salt and hydrogen potassium 
fluoride (HKF 2 ) is an acid salt. 

A solution of hydrofluoric acid and the moist gas both 
attack glass and are used extensively in etching. The glass 
is coated with wax, and the design to be etched is scratched 
through the wax. The glass is then exposed to the gas or 
liquid, which attacks the unprotected places. When the 
wax is removed, a permanent etching like the design is left. 
Glass is an artificial compound of silicon — a silicate. The 
corrosive action of hydrofluoric acid upon glass is due to the 
ease with which the acid decomposes glass and forms with 
silicon a volatile compound, called silicon tetrafluoride (SiF 4 ). 
Equations for the essential chemical changes are : — 

CaSi0 3 + 6 HF = SiF 4 + CaF 2 + 3 H 2 

Calcium Hydrofluoric Silicon Calcium Water 

Silicate Acid Tetrafluoride Fluoride 

Si0 2 + 4HF = SiF 4 + 2 H 2 

Silicon Hydrofluoric Silicon 

Dioxide Acid Tetrafluoride 

Scales on thermometers and on other graduated glass instru- 
ments are often etched with hydrofluoric acid. 

Bromine 

Occurrence. — Bromine is never found free in nature on 
account of its chemical activity. Bromides are widely dis- 
tributed, especially sodium bromide and magnesium bromide. 
The salt springs of Ohio, West Virginia, Pennsylvania, and 




366 INORGANIC CHEMISTRY 

Michigan, and the salt deposits at Stass- 
furt in Germany furnish the main supply 
of the element. Sea water, Chile saltpeter 
(NaN0 3 ), and certain seaweeds contain a 
small quantity of combined bromine. 

Preparation. — Bromine is obtained from 

its native compounds by electrolysis and 

by treatment with chlorine or with sul- 

\y phuric acid and some oxidizing agent like 

Fig. 57.— Appara- potassium chlorate or manganese dioxide. 

bromin^ln^he ^ n ^ e Moratory, bromine is prepared by 

laboratory; Part heating potassium bromide with manganese 

of the liberated dioxide and sulphuric acid in a glass ves- 

romme escapes j (Fig. 57). The bromine is easily liber- 

as a vapor and \ o / j 

part quickly con- ated as a dense, brown vapor, which often 
densestoaliquid condenses to a liquid and runs down the 

which collects in n r j.u i mu u l i 

the V bend of the wa ^ s °* ™ e vessel. The chemical change 
delivery tube. is represented thus : — 

2 KBr + 2 H 2 S0 4 + Mn0 2 = Br 2 + MnS0 4 + K 2 S0 4 + 2 H 2 

Potas- Sulphuric Manga- Bro- Manga- Potas- Water 

sium Acid nese mine nese sium 

Bromide Dioxide Sulphate Sulphate 

Bromine is sometimes prepared in the laboratory by treating 
a bromide with manganese dioxide and hydrochloric acid. 

Bromine is obtained technically from the solution left after 
sodium chloride is crystallized from brine and also from the liquor 
left after potassium chloride is extracted from impure carnallite 
(KC1, MgCl 2 . 6 H 2 0). The hot liquid flows down a large tower 
filled with clay balls ; chlorine gas (frequently obtained by the 
electrolysis of chlorides) and steam forced in at the bottom meet 
the liquid and liberate the bromine, which passes as a vapor out 
of the top into 'a condenser. Equations for the changes are : — 

MgBr 2 + Cl 2 " = Br 2 + MgCl 2 

Magnesium Bromide Chlorine Bromine Magnesium Chloride 

2NaBr + Cl 2 = Br 2 + 2 NaCl 

Sodium Bromide Chlorine Bromine Sodium Chloride 



KClOs 


+ 6HC1 


Potassium 
Chlorate 


Hydrochloric 
Acid 



BROMINE 367 

In the periodic process, used chiefly in the United States, a huge 
stone still is charged with manganese dioxide, hot bittern, and sul- 
phuric acid, and heated by steam. The bromine distills into a con- 
denser, as in the other process. Sometimes potassium chlorate is 
used as the oxidizing agent, and the equation for the essential 
chemical change is : — 

3C1 2 + KC1 + 3H 2 

Chlorine Potassium Water 

Chloride 

Properties. — Bromine is a heavy, reddish brown liquid at 
the ordinary temperature. Its specific gravity is about 3. 
It is a volatile liquid, boiling at about 59° C. The vapor, 
which is given off freely, has a disagreeable, suffocating odor. 
This property suggested the name bromine (from the Greek 
word bromos, a stench). It is poisonous, and burns the flesh 
frightfully. Bromine is somewhat soluble in water. The 
solution called bromine water has a brown color and is some- 
times used instead of bromine itself. The solution contains 
an unstable bromine hydrate. Bromine dissolves in carbon 
disulphide, and this solution is yellow. Its vapor density 
up to about 750° C. requires a molecular weight of 160, and 
since the atomic weight is 79.92, the formula of bromine 
vapor is Br 2 . 

Bromine combines directly with many elements, especially 
hydrogen, phosphorus, and metals. The action is not so 
violent as with chlorine. In fact free chlorine readily dis- 
places bromine from some of its compounds. The chemical 
properties of bromine can be illustrated by a simple experi- 
ment. If powdered magnesium is added to bromine water, 
the brown color disappears, owing to the formation of color- 
less magnesium bromide by the direct combination of mag- 
nesium and bromine. Upon the addition of chlorine (or 
chlorine water) to this colorless solution, the brown color 
reappears, owing to the free bromine which is displaced from 
the magnesium bromide by the more active chlorine. 



368 INORGANIC CHEMISTRY 

Compounds of Bromine. — Hydrogen bromide (HBr) is 
a colorless pungent gas, which fumes in the air and dissolves 
freely in water, forming a solution usually called hydrobromic 
acid. Its other properties closely resemble those of hydro- 
chloric acid, though it is less stable than its chlorine ana- 
logue. Bromides are salts of hydrobromic acid, though 
many are formed by direct combination with bromine. 
Most bromides, like the chlorides, dissolve in water. Po- 
tassium bromide (KBr) is a white solid, made by decom- 
posing iron bromide with potassium carbonate. It is used 
extensively as a medicine and in photography (in preparing 
silver bromide plates and films). 

Miscellaneous. — Bromine itself is used to make potassium 
bromide and other compounds, especially certain coal tar 
dyes. 

Balard discovered bromine in 1826 in the mother liquor 
(or bittern) from brine. Liebig, to whom it was submitted 
for examination, supposed it was chloride of iodine, and thus 
failed to discover its elementary nature, because, as he said, 
he yielded to " explanations not founded on experiment/ ' 

Iodine 

Occurrence. — Free iodine is never found in nature, but 
like chlorine and bromine it occurs in combination with 
metals, especially sodium, potassium, or magnesium. It is 
widely distributed, though the quantity in any one place is 
small. Tobacco, water-cress, cod-liver oil, oysters, and 
sponges contain minute quantities of iodine compounds. 
Native iodides of silver and of mercury are found. The ash 
of some seaweeds contains from 0.5 to 1.5 per cent of its 
weight of iodides of sodium and potassium. Sodium iodate 
(NaI0 3 ) occurs in the deposits of saltpeter in Chile, and is 
now the main source of the element. 



IODINE 369 

Preparation. — Iodine is prepared in the laboratory by a 
method similar to that used for bromine. Potassium iodide, 
manganese dioxide, and sulphuric acid are heated in a glass 
vessel, and the iodine is liberated as a violet vapor, which 
quickly condenses on the upper part of the vessel as dark 
grayish crystals. The equation for the chemical change is — ■ 

2 KI + 2 H 2 S0 4 + Mn0 2 = I 2 + MnS0 4 + K 2 S0 4 + 2 H 2 

Potassium Sulphuric Manganese Iodine Manganese Potassium Water 
Iodide Acid Dioxide Sulphate Sulphate 

On a commercial scale iodine is prepared from the ash of seaweeds 
and from the mother liquors of Chile saltpeter. (1) Along the coasts 
of France, Japan, Scotland, and Norway seaweed is collected and 
burned; sometimes the seaweed is allowed to ferment before being 
burned in order to convert the complex iodine compounds into non- 



Fig. 58. — Apparatus for purifying iodine. 

volatile iodides. From the ash, which is called kelp or varec, the 
soluble portions are removed by agitation with water. The filtered 
liquid is further purified, and from the final mother liquor in which 
the iodides are dissolved the iodine is extracted by heating with 
sulphuric acid and manganese dioxide. Sometimes chlorine is used 
to extract the iodine, the equation for this chemical change being — ■ 



2 Nal 


+ ci 2 


= i 2 + 


2 XaCl 


Sodium 
Iodide 


Chlorine 


Iodine 


Sodium 
Chloride 



In either case the mother liquor and its added ingredients are dis- 
tilled gently in an iron pot with a lead cover, which is connected with 
two rows of bottle-shaped condensers (Fig. 58), The iodine, which 



'370 INORGANIC CHEMISTRY 

collects in these condensers, is purified by washing and resubliming. 
(2) In the other process the mother liquor (which contains about 
22 per cent of sodium iodate) from the Chili saltpeter is mixed 
with sodium sulphite and acid sodium sulphite (HNaSOs); the 
precipitated iodine is collected on coarse cloth, washed, dried, and 
then resublimed, as described above. The equation for this method 
of preparation is — 

2NaI0 3 + 3Na 2 S0 3 + 2 HNaS0 3 = I 2 + 5Na 2 S0 4 4- H 2 

Sodium Sodium Acid Sodium Iodine Sodium Water 

Iodate Sulphite Sulphite Sulphate 

Courtois, a French chemist, discovered iodine in 1812, in an attempt 
to prepare potassium nitrate from seaweed. Davy and Gay-Lussac 
established its elementary nature and discovered many of its proper- 
ties. The present name was given by Davy. 

Properties. — Iodine is a dark grayish crystalline solid, 
resembling graphite in luster. It crystallizes in plates which 
have a specific gravity of about 5. It is volatile at the ordi- 
nary temperature, melts at 114° C, and boils at 184° C. 
When gently heated, the vapor which is formed has a beauti- 
ful violet color. This color suggested the name of iodine 
(from the Greek word iodes, violetlike). The vapor is nearly 
nine times heavier than air. It has an odor resembling that 
of dilute chlorine, though less irritating. When the vapor 
is heated, its color changes from violet to deep blue, and the 
density decreases. Experiment shows that from about 
200 to 700° C. the vapor density requires a molecular weight 
of about 255; since the atomic weight is 126.92, the mole- 
cules contain only two atoms and the formula of iodine 
vapor is I 2 up to this temperature. As the temperature 
rises the molecules dissociate, until at about 1700° C, the 
vapor consists entirely of atoms. Iodine stains the skin 
yellow, and turns cold starch solution blue. The presence 
of a minute trace of iodine may be thus detected, one part 
of iodine in over 400,000 parts of water producing the blue 
color. The exact nature of this blue substance is unknown. 



IODINE 371 

The presence of starch in many vegetable substances can be 
shown by this delicate test. Iodine dissolves slightly in 
water, and freely in -alcohol, chloroform, carbon disulphide, 
and potassium iodide solution. The chloroform and carbon 
disulphide solutions are violet, but the others are brown or 
even black. 

The chemical properties of iodine resemble those of chlorine 
and bromine, but it is less active. It forms no hydrate with 
water, differing from chlorine and bromine in this respect. 
Bromine and chlorine displace iodine from its binary com- 
pounds, chlorine and chlorine water being often used for this 
purpose. It combines directly with some non-metals and 
most metals. 

Compounds of Iodine resemble the corresponding ones of 
chlorine and bromine. Hydriodic acid (HI) is much like 
hydrobromic and hydrochloric. Iodides are salts of hy- 
driodic acid, and like many salts they are prepared in various 
ways. In general behavior they are similar to bromides and 
chlorides. Potassium iodide (KI) is prepared and used like 
potassium bromide. 

Brief reference was made in Chapter XI to the equilibrium 
established when hydriodic acid is formed by heating hydro- 
gen and iodine in a closed tube. The elements begin to 
combine at about 445° C, while the acid begins to decom- 
pose at about 180° C. The reactions proceed in direct 
and reverse directions, though not to completion, for a con- 
dition of equilibrium is soon reached and further accumula- 
tion of the products ceases. Chemical change is still going 
on, but one reaction neutralizes the other. For example, at 
about 445° C. a condition of equilibrium prevails, both reac- 
tions are still in operation, though the tube contains about 
79 per cent of hydriodic acid and 21 per cent of hydrogen 
and iodine. This state of equilibrium is maintained as long 



372 INORGANIC CHEMISTRY 

as certain conditions prevail, but like other states of equilib- 
rium it is rather sensitive and is easily displaced by varying 
the temperature, the pressure, or the concentration of the 
substances in the tube. The equation expressing the revers- 
ible reaction just discussed may be written — 

H 2 + I 2 ^I2HI 

Miscellaneous. — Iodine dissolved in alcohol, in potassium 
iodide solution, or both, is used (under the name of tincture 
of iodine) as an application for the skin to prevent the spread 
of eruptions or to reduce swellings. Iodine is used to make 
medicinal preparations, especially iodoform (CHI 3 ), which 
is used as an antiseptic dressing for wounds. Large quan- 
tities of iodine are used in making aniline dyes. 

Halogen Family. — The physical and chemical properties 
of the halogen elements furnish a typical illustration of the 
resemblances, differences, and gradation of properties which 
characterize a family of elements in the same periodic group- 
If these elements are arranged in the order of their atomic 
weights, from fluorine (19.0) through chlorine (35.46) and 
bromine (79.92) to iodine (126.92), the periodic nature of 
the group is revealed. Thus, the specific gravity increases 
in this order {i.e. fluorine to iodine) ; the color likewise grows 
deeper, but the volatility decreases. So also, the melting 
points of the solidified elements and the boiling points of 
the liquefied elements increase in this order. Chemically, 
these elements unite with hydrogen to form analogous com- 
pounds, and the intensity of the chemical action decreases 
gradually as we pass from fluorine to iodine. 

The halogen acids resemble each other in their solubility, 
all (except HF) forming solutions in which the dissociation 
is large; in other words, they are strong acids. Salts of the 



IODINE 373 

halogen acids often resemble one another, especially those 
of sodium and potassium, which are white solids, soluble in 
water, from which they crystallize in cubes. 

The valence of the halogens is one toward hydrogen and 
metals. 

Problems and Exercises 

1. What is the percentage composition of (a) fluor spar (CaF 2 ) 
and (6) cryolite (Xa 3 AlF 6 )? 

2. How much (a) calcium sulphate and (b) hydrofluoric acid are 
formed by heating 100 gm. of fluor spar with sulphuric acid? 

3. Calculate the percentage composition of (a) potassium bromide 
(KBr), (6) potassium iodide (KI), (c) silver bromide (AgBr), and 
(d) iodoform (CHI 3 ). 

4. How much potassium iodide is needed to prepare 63.5 gm. of 
iodine? 

5. How much potassium bromide is needed to prepare 10 gm. of 
bromine ? 

6. If 3.946 gm. of silver are needed to precipitate the bromine in 
4.353 gm. of potassium bromide, what is the atomic weight of bro- 
mine? 

7. If the specific gravity of bromine is 3, what volume does one 
pound occupy? 

8. If 63.8351 gm. of silver iodide yield 38.9656 gm. of silver chloride, 
what is the atomic weight of iodine if 107.88 and 35.46 are accepted 
as the atomic weights of silver and chlorine respectively? 

9. Write the formulas of the following compounds by applying 
the principle of valence (see Chapter XIV) : Lithium chloride, mag- 
nesium iodide, mercurous iodide, mercuric bromide, aurous chloride, 
auric bromide, barium fluoride, silicon fluoride, platinum chloride, 
zinc chloride, ferrous bromide, ferric chloride, lead iodide. 

10. Write the equations for the following reactions: (a) Potassium 
bromide and silver nitrate form silver bromide and potassium nitrate. 
(6) Sodium iodide and silver nitrate form silver iodide and sodium 
nitrate. Write the corresponding ionic equations. 

11. Calculate the simplest formulas : (a) F = 48.72, Ca = 51.28 ; 
(b) Br = 67.22, K = 32.77; (c) I = 76.5, K = 23.49. 

12. The formulas of bromic acid and iodic acid are HBr0 3 and 
HI0 3 respectively. Write the formulas of their salts corresponding 
to K, calcium, Mg, sodium, and Ba. 



CHAPTER XXI 
Boron 

Occurrence. — Boron (B) is never found free, but several 
of its compounds are abundant, eg. borax (Na 2 B 4 7 ), boric 
acid (H3BO3), boracite ((Mg 3 B 8 15 ) 2 . MgCl 2 ), and colemanite 
(Ca 2 B 6 O u .5H 2 0). 

Boron is prepared by heating the oxide (B 2 3 ) with mag- 
nesium or the chloride (BCI3) with hydrogen. It is black, 
non-crystalline, and very hard. It unites with oxygen and 
with nitrogen, forming the oxide (B 2 3 ) and a nitride (BN). 
When heated in the electric furnace with carbon, it forms 
carbon boride (CB 6 ), which is nearly as hard as diamond. 

Boric Acid, H 3 B0 3 , is contained in the waters and steam of 
certain volcanic regions, notably Tuscany in Italy. Large 
basins or tanks are built around these steam jets (called 
suffioni) and are arranged so that the water flows at intervals 
from one reservoir into the next lower, constantly becoming 
charged with more boric acid, as the steam condenses. The 
final solution is evaporated by aid of the heat from the steam 
jets, and the crude boric acid which settles out is purified 
by recrystallization. This compound is sometimes called 
boracic acid. Boric acid is made in the United States from 
borax and colemanite, and in Germany from the boracite 
found at Stassfurt. The essential feature of the process is a 
decomposition of the mineral into boric acid by an inorganic 
acid. The following equation illustrates the transforma- 
tion : — 



fa 2 BA + H-jSOi + 5H 2 


= 4H 3 B0 3 


+ Na 2 SO« 


Borax Sulphuric 


Boric 


Sodium 


Acid 


Acid 


Sulphate 



374 



BORON 375 

Boric acid crystallizes in lustrous, white flakes, which fee) 
greasy. It dissolves slightly in cold water, readily in hot 
water and in alcohol. When the alcoholic solution is burned, 
a boron compound colors the vapor green. This is a test 
for boron compounds. 

Boric acid is used in making borax, in the manufacture of 
enamels and glazes for pottery, as an antiseptic in medicine 
and surgery, and for preserving meat, fish, milk, butter, beer, 
and wine. 

When boric acid is heated, it loses water and is transformed 
into metaboric acid (HB0 2 ) at 100° G. and tetraboric acid 
(H 2 B 4 7 ) at 140° C. Boric acid forms no salts, but the other 
acids do, the best-known salt being sodium tetraborate or 
borax. 

Borax, Na 2 B 4 7 , occurs native in California, and an impure 
borax called tinkal comes from the East. Most of the com- 
mercial borax is made from calcium borate (colemanite) by 
boiling with sodium carbonate thus : — 

2 Ca 2 B 6 O n +4 Na 2 C0 3 +H 2 = 3 Na 2 B 4 7 +4 CaC0 3 +2 NaOH 

Calcium Sodium Borax Calcium Sodium 

Borate Carbonate Carbonate Hydroxide 

Borax is a white crystalline solid and has five or ten mole- 
cules of water of crystallization, depending upon the tem- 
perature at which crystallization occurs. It effloresces in 



3 ^O 



Fig. 59. — Looped platinum wire for making tests with borax beads. 

the air. When crystallized borax is heated, it swells up (as 
the water of crystallization escapes) into a white porous 
mass, which finally becomes a glassy solid. If the borax 
is melted on the end of a looped platinum wire, the trans- 
parent globule is called a borax bead (Fig. 59). This 



376 



INORGANIC CHEMISTRY 



glassy borax dissolves metallic substances, especially metallic 
oxides. These beads may be made to assume different colors 
characteristic of the metals, when the beads are heated with 
the oxides or solutions of different metals. The following 
table shows the — 

Colors of Borax Beads 



Metal 


Oxidizing Flame 


Reducing Flame 




Hot 


Cold 


Hot 


Cold 


Chromium 
Cobalt 
Copper 
Manganese 


Reddish yellow 

Blue 

Green 

Violet 


Yellowish green 
Blue 

Greenish blue 
Violet 


Green 
Blue 
Colorless 
Colorless 


Green 
Blue 
Red 
Colorless 



The bead test is often used in chemistry to confirm other 
observations or to suggest further examination. The chem- 
ical changes in borax beads can be readily understood if 
borax is regarded as a mixture of sodium metaborate and 
boron trioxide (2NaB0 2 + B 2 3 ); the acid oxide (B 2 3 ) 
unites with the basic {i.e. metallic) oxides and forms a colored 
borate. For example, with copper oxide a green bead is 
obtained which is 2 NaB0 2 , Cu(B0 2 ) 2 . 

A solution of borax has an alkaline reaction, because borax 
hydrolyzes with water and boric acid is only slightly ionized. 

Large quantities of borax are used in the manufacture of 
enamels and glazes, especially those which form the protective 
coating of domestic utensils. Considerable is used for 
preserving canned meat and fish. It is a cleansing agent, 
and large quantities are consumed in laundries as well as in 
the manufacture of soap, particularly the variety intended 
for use in hard water. (See Soap and Hard Water.) The 
property of dissolving oxides adapts borax for use in soldering 
certain metals. Solder adheres only to clean metals, so a 



BORON 377 

little borax is used to dissolve the film of oxide on the surfaces 
to be joined. It also finds use as a mordant in calico printing 
and in dyeing, and in the manufacture of water-soluble 
varnishes. It is an ingredient of the ointments, lotions, and 
powders which are designed to relieve hoarseness or skin 
eruption. 

Miscellaneous. — Boron is a non-metal and belongs to the 
aluminium family in the third periodic group (III), but chem- 
ically it closety resembles carbon and silicon and their com- 
pounds. Its valence is three. 

Problems 

1. What per cent of borax (Xa 2 B 4 7 . 10 H 2 0) is boron? 

2. Write the equations for the transformation of boric acid into 
metaboric and tetraboric acids; and the equation for the formation 
of boron trioxide from tetraboric acid. 

3. How many grams of H3BO3 are formed by the interaction of 
water, sulphuric acid, and 782 gm. of borax? 

4. Calculate the per cent of boron in (a) colemanite, and (b) bo- 
racite. 

5. How many grams of borax can be made from a metric ton of 
colemanite ? 

6. Write the formulas of the barium and aluminium salts of 
tetraboric and metaboric acids and calculate the per cent of boron 
in each salt. 



CHAPTER XXII 
Silicon — Glass 

Occurrence. — Silicon does not occur free in nature, being 
found almost exclusively as silicon dioxide (Si0 2 ) or as sili- 
cates. These compounds are so abundant and widely dis- 
tributed that approximately one fourth of the earth's crust 
is silicon. Sand and the different varieties of quartz are silicon 
dioxide, while many rocks are silicates. 

Preparation and Properties. — Silicon is no longer a rare 
element. It is prepared by heating a special mixture of 
silicon dioxide and carbon in an electric furnace at a temper- 
ature which is carefully regulated to prevent loss of the silicon 
by volatilization or combination with the carbon. The 
equation for the chemical change is — 

Si0 2 + 2C = Si + 2 CO 

Silicon Carbon Silicon Carbon 

Dioxide Monoxide 

Thus prepared silicon is a gray-black, lustrous, brittle 
solid. It melts at about 1400° C. and oxidizes to silicon 
dioxide when heated to about this temperature. The spe- 
cific gravity is about 2.37. It is almost as hard as quartz. 
At high temperatures silicon and oxygen form silicon dioxide ; 
with certain elements silicon forms silicides. Silicon and the 
halogens form volatile compounds, e.g. silicon tetrafluoride 
(SiF 4 ) (see pages 365, 384). With sodium hydroxide it 
forms sodium silicate (Na 4 Si0 4 ) and hydrogen. 

Silicon Dioxide or Silica, Si0 2 , is the most common com- 
pound of silicon. It occurs native in both crystalline and 

378 



SILICON 



379 




amorphous conditions, but when produced by a chemical 
process in the laboratory it is usually a white amorphous 
powder. Sand, gravel, sandstone, and quartzite are almost 
wholly silica. It is an essential ingredient of many rocks, 
as granite and gneiss. Quartz is silicon dioxide. Pure crys- 
talline quartz is colorless and transparent, and is frequently 
found as crystals which consist usually of a six-sided prism 
with a six-sided pyramid at one or both ends ; but the crystals 
are often distorted or complex (Fig. 60). There are many 
varieties of quartz, which differ 
in color and structure, due to 
minute impurities or to the mode 
of formation. Among the crys- 
talline varieties are the clear, 
colorless rock crystal, the purple 
amethyst, and the rose, yellow, 
glassy, milky, and smoky forms. 
Varieties imperfectly crystalline 
or amorphous are the waxlike chalcedony, the various forms 
of agate having different colored layers, the reddish brown 
carnelian, the black and white onyx, the red or brown jasper, 
the dull brown or black flint, and the brittle chert. Opal is 
hydrated silica (Si0 2 .nH 2 0). Petrified or silicified wood 
is largely some variety of quartz which has replaced the 
woody fiber. Infusorial or diatomaceous earth is a variety 
of silica consisting of the shells of minute organisms called 
diatoms. 

Quartz crystals and most crystalline varieties of silica 
are hard enough to scratch glass. They are insoluble in 
water and acids, except hydrofluoric acid, but are transformed 
into a soluble alkaline silicate when heated in the hydrox- 
ides or carbonates of sodium and potassium. Thus, when 
fine sand is fused with sodium carbonate the equation for 
the reaction is — 



Fig. 60. — Quartz crystals. 



380 INORGANIC CHEMISTRY 



Na 2 C0 3 


+ 


Si0 2 


Na 2 Si0 3 


+ 


co 2 


Sodium 




Silicon 


Sodium 




Carbon 


Carbonate 




Dioxide 


Silicate 




Dioxide 



Silica itself is infusible, except in the oxyhydrogen flame and 
electric furnace. If pure silica is fused with certain pre- 
cautions, the molten mass can be shaped into elastic threads, 
which are used to suspend delicate parts of electrical instru- 
ments, and into tubes, flasks, crucibles, etc., which do not 
crack by sudden heating and cooling. The specific gravity 
of quartz is about 2.65. 

Sandstone and quartzite are used as building stones, and 
hard sandstone is made into grindstones and whetstones. 
Sand is used in making sandpaper, glass, porcelain, and 
mortar. Glass is roughened and cut by blowing or " blast- 
ing " fine sand against it. Many of the varieties of quartz 
are cut and polished into ornaments and gems, e.g. amethyst, 
opal, and agate. Rock crystal is used as the " diamond " 
in cheap jewelry, and is sometimes cut into lenses for eye- 
glasses and optical instruments. Petrified wood is cut and 
polished into table tops, mantelpieces, and fireplaces. In- 
fusorial earth is used to polish silver (" electro-silicon " 
being the commercial name of one kind) and in making 
cement, " soluble glass/ 7 dynamite, and refractory brick. 

Silica and Plants. — The ash of many plants contains silica, 
showing that some compound of silicon is assimilated by 
the plant from the soil — probably silicic acid or a soluble 
silicate (see below). The ash of rye and wheat straws and 
of potato stems contains from 40 to 70 per cent of silica. 
Plants like horsetail and bamboo are rich in silica. The 
silica is probably not a plant food in the strict sense, but 
gives firmness to the tall stalks, especially to their joints. 

Silicic Acids and Silicates. — Silicon being a non-metal 
forms acids, many of which are complex and known only 



SILICON 381 

through the corresponding salts. The two which are simple 
and best known are metasilicic acid (H 2 Si0 3 ) and orthosilicic 
acid (H 4 Si0 4 ). As stated in a preceding paragraph, sodium 
silicate (Na 2 Si0 3 ) is formed when silicon dioxide is fused 
with sodium carbonate. Now sodium silicate dissolves in 
water, and when hydrochloric acid is added to a concentrated 
solution, a silicic acid is precipitated as a white gelatinous 
mass. The precipitate is probably orthosilicic acid, but 
when this precipitate is dried only metasilicic acid is found 
in the residue, thus : — 



H 4 Si0 4 = 


H 2 Si0 3 + H 2 


Orthosilicic 


Metasilicic Water 


Acid 


Acid 



The metasilicic acid on further heating decomposes into 
silicon dioxide and water, thus : — 



H 2 SiO s 


= Si0 2 + H 2 


Metasilicic 


Silicon Water 


Acid 


Dioxide 



It appears then that these two silicic acids are closely related. 
Indeed orthosilicic acid may be regarded as a sort of parent 
of the other silicic acids, which, although they have not been 
isolated, may be conveniently thought of as molecules of 
orthosilicic acid minus one or more molecules of water. 
Following out this relationship we have for example the fol- 
lowing hypothetical silicic acids : — 

2 H 4 Si0 4 - H 2 = H 6 Si 2 7 , or Disilicic Acid 

3 H 4 Si0 4 -4 H 2 = H 4 Si 3 8 , or Trisilicic Acid 

Colloidal Silicic Acid. — Sodium silicate and hydrochloric 
acid do not always interact as described above. If the 
sodium silicate solution is dilute, or the hydrochloric acid 
concentrated or in excess, then the silicic acid which is 



382 INORGANIC CHEMISTRY 

formed remains in solution as colloidal silicic acid. It can- 
not be filtered out, though it can be separated by dialysis 
from the sodium chloride in the solution. Thus, if the col- 
loidal solution of silicic acid is placed in a vessel having a 
bottom of parchment and hanging in a larger receptacle 
filled with water, the silicic acid will be retained in the 
smaller vessel but the sodium chloride will pass through the 
parchment into the water. This process was devised by 
Graham, who did the first work on colloids, ^a^ ujXlJ l^CoX^^ 

Colloidal Solutions. — In a colloidal solution the substance is 
suspended as exceedingly fine particles, which will pass through 
filter paper but not through parchment or other animal membranes ; 
neither will the particles settle, though they can be precipitated (or 
coagulated) under special conditions. Substances which form col- 
lodial solutions are called colloids, and while in solution are in the 
colloidal condition. Colloidal solutions are not solutions in the 
usual sense, for unlike true solutions they show very little, if any, 
elevation of the boiling point or depression of the freezing point ; 
moreover, when a converging beam of strong light is passed through 
a colloidal solution, the path is brightened by the particles, whereas 
a true solution remains dark. Sometimes colloidal solutions are 
called colloidal suspensions. 

The term colloid (from the Greek word for glue), which was 
first applied to sticky substances, now includes two general classes 
of substances : (1) Those, like, agar-agar and gelatin, that form 
jellylike masses on cooling or concentration ; and (2) those, like gold 
and arsenious sulphide, that coagulate {i.e. precipitate upon the 
addition of an electrolyte). Certain colloids are precipitated merely 
by heating. Many colloids carry electrical charges. Arsenious 
sulphide (As 2 S 3 ), silver chloride (AgCl), and certain metals (Ag, 
Cu, Pt) are negative, while ferric hydroxide (Fe(OH) 3 ) and many 
basic substances are positive ; starch and gelatin are neutral. 
Charged colloids are precipitated by oppositely charged ions and 
colloids. Thus, colloidal arsenious sulphide is precipitated by hy- 
drogen ion {e.g. from hydrochloric acid) and also by colloidal ferric 
hydroxide ; so also negative metaphosphoric acid and positive 
(usually) albumin precipitate one another (see page 399). 

Coagulation of colloids can be retarded or prevented by adding 



SILICON 383 

a protective colloid. Thus, gelatin is often used to render col- 
loidal silver bromide more stable, e.g. in photographic plates. It 
is supposed that the protective colloid forms a film around the 
other colloid and thereby prevents diffusion. 

Silicates are the salts of silicic acids, though their cor- 
responding acids have not been isolated in most cases. So- 
dium and potassium silicates are salts of the well-known 
metasilicic acid (H 2 Si0 3 ) . They are the only silicates soluble 
in pure water, and the thick, sirupy solution of each (or 
both) is called water glass. It finds extensive use in the 
manufacture of soap, certain cements, artificial stone, and 
fireproof materials. As already stated many rocks and min- 
erals are silicates and make up a large part of the earth's 
crust. The following list shows the relations of a few 
silicates to their acids : — 

Metasilicates I Wollastonite > CaSi0 3 

/a u c TjQ-rw i Enstatite, Mgfei0 3 
(Salts of H 2 Si0 3 ) -d , -d A1 %-^ v 
[ Beryl, Be 3 A] 2 (Si0 3 )6 

Zircon, ZrSi0 4 
Kaolin, H 2 AJ 2 (Si0 4 ) 2 . H 2 
L Olivine, Mg 2 Si0 4 . Fe 2 Si0 4 

Disilicate Serpentine, Mg 3 Si 2 7 . 2 H 2 

(Salt of H 6 Si 2 7 ) 

Trisilicate Orthoclase, KAlSi 3 8 

(Salt of H 4 Si 3 8 ) ' 

Other silicates are mica, hornblende, augite, slate, talc, lava, 
feldspars related to orthoclase, asbestos, garnet, and tour- 
maline. The most abundant are the silicates of calcium, 
aluminium, magnesium, potassium, sodium, and iron. 

Silicic acid is a feeble acid. It does not redden blue litmus 
nor liberate hydrogen when added to magnesium. Never- 
theless it is properly called an acid because it forms salts. 
These salts, which we have already seen make up a large part 



Orthosilicates 
(Salts of H 4 Si0 4 ) 



384 INORGANIC CHEMISTRY 

of the earths crust, are slowly decomposed by carbonic acid, 
i.e. by the joint action of the carbon dioxide and water vapor 
in the atmosphere. This disintegration of the silicates is 
called weathering. 

The water of many hot springs, as in the Yellowstone 
Park, contains alkaline silicates ; and when the solution comes 
to the surface, some of the silicate is decomposed by the 
carbon dioxide in the air, and the silica is deposited around 
the spring in beautiful forms called geyserite or siliceous 
sinter. 

Silicon Tetrafluoride, SiF 4 , is a colorless gas which has a 
pungent, suffocating odor. It is formed when hydrofluoric 
acid interacts with silicon dioxide or silicates. Thus, with 
silicon dioxide the equation is — 

Si0 2 + 4HF = SiF 4 + 2H 2 

Silicon Hydrofluoric Silicon 

Dioxide Acid Tetrafluoride 

Silicon tetrafluoride forms fumes in moist air and interacts 
readily with water, thus: — 



3 SiF 4 + 


4H 2 = 


= H 4 Si0 4 


+ 2 H 2 SiF 6 


Silicon 
Tetrafluoride 




Silicic 
Acid 


Hydrofluosilicic 
Acid 



The hydrofluosilicic acid (sometimes called simply fluosilicic 
acid) remains in solution, while the silicic acid is precipitated. 
The formation of the white gelatinous silicic acid when the 
gases from the interaction of hydrofluoric acid and a com- 
pound of silicon are led into water is often used as a test for 
silicon. 

Siloxicon is the commercial name of a highly refractory 
substance produced by heating a mixture of silicon dioxide 
and carbon to about 2500° C. in a special form of an electric 



SILICON 385 

furnace. It is not a definite compound, but varies in com- 
position from Si 2 C 2 to Si 7 C 7 0. It is a gray-green, granular 
powder which can be readily shaped into bricks, linings, and 
other forms of refractory articles. 

Silicides are compounds of silicon and other elements. 
Carborundum or carbon silicide is the most important. 

Carborundum, SiC, is a crystalline compound consisting 
solely of silicon and carbon. It varies in color from white 
to emerald green and is sometimes iridescent. It is ex- 
tremely hard, being nearly as hard as diamond. The specific 
gravity is about 3. Acids do not affect it, but it is decom- 
posed by fusing it with potassium hydroxide and other alkalies. 

Its extreme hardness has led to its application as an abra- 
sive, and large quantities are made into a great variety of 
grinding wheels, whetstones, and polishing cloths. 

Carborundum is manufactured by fusing a mixture of 
sand and coke in an electric furnace constructed on the 
resistance type (Fig. 61). It is essentially an oblong box 
with permanent ends and loosely built sides. Each end is 
provided with a heavy metal plate. The wires for the 
electric current are attached to the outer ends of these 
plates, while the huge carbon electrodes fit into the inner 
ends, and project into the furnace. A cylindrical mass of 
granulated coke makes an electrical connection between the 
electrodes. The mixture of sand and coke (to which salt and 
sawdust are added to contribute to the fusion and porosity) 
is packed around this core inside the box. The heat gen- 
erated by the resistance of the carbon core to the passage of 
the powerful current of electricity produces a chemical 
change essentially as follows : — 

Si0 2 + 3C = SiC + 2 CO 

Sand Carbon Carborundum Carbon 

Monoxide 



386 



INORGANIC CHEMISTRY 




•3 

Pi 

u 
O 

•s 






The change is due solely to the intense heat, i.e. it is an 
electrothermal, not an electrolytic change. When the opera- 
tion is over and the furnace is cool, the side walls are pulled 
down, and the carborundum i& removed. The purest grade 



GLASS 387 

is found around the core. The product is crushed, treated 
with sulphuric acid to remove the impurities, washed, dried, 
and graded according to the size of the particles. 

Miscellaneous. — Silicon is a non-metallic element and 
belongs to the carbon family in the fourth (IV) periodic 
group. Certain physical properties suggest a metallic char- 
acter, but chemically silicon is very closely related to carbon. 
Both have allotropic modifications and form analogous com- 
pounds, e.g. C0 2 and Si0 2 , CH 4 and SiH 4 . Both form many 
compounds of great importance, so that w r e might conven- 
iently regard silicon as the chief element in the mineral king- 
dom, just as carbon is in the organic realm. 

Silicon has the valence of four in its compounds. 

Glass 

Glass is an amorphous, more or less transparent solid. It 
is a homogeneous mixture of silicates with an excess of 
silica. Glass is not made by mixing silicates, but by fusing 
a mixture of sand, an alkali, and a calcium or lead compound. 
The alkali is potassium carbonate (K 2 C0 3 ) or sodium car- 
bonate (Na 2 C0 3 ), though sodium sulphate and sodium nitrate 
are used in some cases as auxiliary substances; the calcium 
compound is limestone (CaC0 3 ) or lime (CaO) ; and the lead 
compound is litharge (PbO) or red lead (Pb 3 4 ). Besides 
these fundamental ingredients, small quantities of other 
substances are used, e.g. (1) broken glass (called cullet), 
which lowers the melting point of the mixture ; (2) arsenic 
trioxide (As 2 3 ), which destroys carbonaceous impurities; 
(3) carbon, which lowers the melting point when sodium 
sulphate is used and likewise imparts a color from straw to 
amber ; and (4) manganese dioxide (Mn0 2 ) , which neutralizes 
the green color caused by iron compounds (often present in 
impure materials). 



388 INORGANIC CHEMISTRY 

The process consists in melting a carefully prepared mix- 
ture of the proper ingredients in a refractory fire-clay pot. 
The heat is often obtained by burning gas — manufactured 
or natural. During the melting, gases escape, and the im- 
purities, which rise to the surface as a scum, are removed. 
The molten mass is allowed to cool until it has the proper 
consistency. A portion is then collected as a soft ball on 
the end of an iron tube and brought to the desired shape, 
either by forcing it into a mold or by blowing into the tube 
and simultaneously manipulating the plastic mass by twist- 
ing and swinging. The details of the procedure, however, 
vary with the article being made. Many objects, such as 
tumblers and small dishes, are now made by pressing the 
plastic glass with a die or by blowing it into a mold. 
Fruit jars, bottles, and lamp chimneys are blown by 
machinery. 

All glass must be cooled slowly to prevent brittleness. 
This operation is called annealing, and is accomplished by 
passing the objects slowly through a furnace in which the 
temperature is gradually lowered. 

There are four possible kinds of glass and many varieties 
of each. Their properties depend upon the proportion of the 
ingredients, and each kind may be made to approach the 
others in properties by varying these proportions. Arranged 
in the order of their fusibility and beginning with the soft- 
est, the four kinds of glass are: (1) Sodium-lead glass, (2) 
potassium-lead glass, (3) sodium-calcium glass, and (4) 
potassium-calcium glass. Flint glass is a lead glass; it is 
lustrous, refracts light to a high degree, and is made into 
ornaments, lenses for optical instruments, and also into 
shades for electric and gas lights. Cut glass objects are 
made from flint glass by means of simple grinding and pol- 
ishing machinery. Window, plate, crown, table, and bottle 
glass is a sodium-calcium glass ; it is sometimes called 



GLASS 389 

soda glass or soft glass to distinguish it from the potassium- 
calcium glass, which is hard. 

Window glass is. made by blowing a lump of glass into a 
hollow globe and then into a cylinder; this, on being opened 
at both ends and cut lengthwise, spreads out flat. Plate 
glass is made by pouring the molten glass upon a large table, 
rolling it with a hot iron roller, and subsequently grinding and 
polishing it until the surfaces are parallel. Plate glass is 
used for large windows and for mirrors, but considerable 
rough plate is used for skylights and floors. Crown glass is 
a superior quality of window glass. It has a brilliant sur- 
face and is used as "bull's-eyes" in decorative windows and 
as lenses for optical instruments (in conjunction with flint 
glass). Most chemical glassware is sodium-calcium glass. 
It therefore softens when heated and the flame becomes 
yellow from the sodium. Bohemian or hard glass is a potas- 
sium-calcium glass. It is much harder than the other kinds 
and is used in making chemical apparatus designed to with- 
stand great heat. Soft glass is slightly soluble in water, but 
hard glass is less so, hence special varieties of hard glass are 
often made into apparatus which resists the solvent action 
of water and chemical reagents; Jena glass is one variety. 

Colored glass is made by adding different substances to 
the mixture. Iron, chromium, and certain copper com- 
pounds make it green, the green color of many bottles and 
fruit jars being due to iron compounds in the impure ma- 
terials used ; cobalt compounds produce different shades of 
blue ; manganese dioxide gives a pink or a violet color ; 
yellow is produced by charcoal, sulphur, uranium compounds, 
or silver ; the deep red glass so extensively used in lanterns 
is usually colored by selenium compounds ; milky glass is 
made by adding calcium phosphate, fluor spar, or cryolite ; 
stained glass is ordinary glass to which fusible pigments are 
applied with a brush and then fixed by heat ; iridescent glass 



390 INORGANIC CHEMISTRY 

is made by secret processes, though it is known that one 
consists in exposing a special variety of absorbent glass to 
the vapors of metallic oxides. 



Problems and Exercises 

1. How can Si0 2 be transformed into H 2 Si0 3 ? How many 
grams of Si0 2 are needed for 75 gm. of H 2 Si0 3 ? 

2. How can Si0 2 be transformed into H 2 SiF 6 ? How much Si0 2 
is needed for 100 gm. of H 2 SiF 6 ? 

3. How much hydrofluosilicic acid can be made by the inter- 
action of water and a metric ton of silicon tetrafluoride ? 

4. Calculate the atomic weight of silicon from the following 
data : (a) 2.621 gm. of silicon tetrachloride (SiCl 4 ) required 6.6445 
gm. of silver for precipitation of the chlorine ; the atomic weight 
of silver was accepted as 107.88. (b) 95.52367 gm. of silicon tetra- 
bromide (SiBr 4 ) yielded 16.56868 gm. of silicon dioxide ; the atomic 
weights of oxygen and bromine were accepted as 16 and 79.92 
respectively. 

5. Write the formulas of the following compounds by applying 
the principle of valence (see Chapter XIV) : Silicon iodide, hydrogen 
silicide, silicon sulphide, carbon silicide. 

6. Calculate the per cent of silicon in (a) calamine, Zn 2 Si04, 
(&) chrysocolla, CuSi0 3 , (c) analcite, Na 2 Al 2 Si 4 O i2 , (d) heulandite, 
CaAl 2 Si 6 0i 6 , (e) beryl, Be 3 Al 2 Si 6 0i 8 (Be = 9), (/) phenacite, Be 2 Si0 4 , 
(g) garnet, Ca 3 Al 2 Si 3 0i 2 , (h) muscovite, KAlSi0 4 , (i) olivine, 
Mg 2 Si0 4 . FeS0 4 , (j) zircon, ZrSi0 4 (Zr = 90.6). 

7. How much silicon can be made (a) by reducing a metric ton 
of sand (90 per cent pure) with C? (b) from 119 gm. of potassium 
silico-fluoride? (Equation is K 2 SiF 6 + 4 K = Si + 6 KF.) 

8. How much metasilicic acid (H 2 Si0 3 ) can be made by the in- 
teraction of water and a metric ton of silicon tetrafluoride ? 

9. Suppose opal is Si0 2 . 10 H 2 ; calculate its per cent of (a) 
Si0 2 , (b) Si, (c) H 2 0. 

10. (a) How much metasilicic acid can be made from a metric 
ton of orthosilicic acid? (b) How much Si0 2 from a metric ton of 
metasilicic acid? 

11. Calculate the simplest formulas corresponding to (a) Si — 
35.897, H = 2.564, = 61.538; (b) Si = 29.166, H=4.166, = 66.666. 
What is the name of each compound? 



CHAPTER XXIII 
Phosphorus, Arsenic, Antimony, and Bismuth 

Phosphorus, arsenic, antimony, and bismuth, together 
with nitrogen, belong to the nitrogen family in the fifth (V) 
periodic group of elements. 

Phosphorus 

Occurrence. — Free phosphorus is not found in nature. 
But phosphates are numerous and abundant, the most 
common being phosphorite (" phosphate rock/' impure 
Ca 3 (P0 4 ) 2 ) and apatite (Ca 5 F(P0 4 )3). Approximately .1 
per cent of the earth's crust is phosphorus. Phosphates are 
present in fertile soils and some iron ores. Phosphorus 
compounds are essential constituents of seeds, and also of 
the brain, nerves, muscles, and bones of animals. 

Phosphorus was discovered in 1669 by Brand, a German alchemist, 
who obtained it by heating a certain kind of animal matter. Scheele, 
in 1771, extracted it from bones. 

Preparation. — Phosphorus is prepared industrially by 
two processes: (1) In the older process bone ash (which is 
over 80 per cent calcium phosphate) or a native phosphate is 
finely ground and mixed in large vats w T ith enough sulphuric 
acid to produce the following change : — 

Ca 3 (P0 4 ) 2 + 3 H 2 S0 4 = 2 H 3 P0 4 + 3 CaS0 4 

Calcium Sulphuric Phosphoric Acid Calcium 

Phosphate Acid (Ortho-) Sulphate 

The insoluble calcium sulphate is removed by filtering the 
mixture through cinders. The phosphoric acid solution is 

391 



392 



INORGANIC CHEMISTRY 



concentrated, mixed with sawdust, coke, or charcoal, and 
dried, being changed thereby into metaphosphoric acid ac- 
cording to the equation : — 

H3PO4 = HPO3 + H 2 



Phosphoric Acid 
(Ortho-) 



Phosphoric Acid 
(Meta-) 



The dried mass is heated to a high temperature in clay 
retorts arranged in tiers (Fig. 62), the change thus produced 
being substantially — 

+ 12C - P 4 + 2H 2 

Carbon Phosphorus Hydrogen 



4HP0 3 

Phosphoric Acid 
(Meta-) 



+ 12 CO 

Carbon 
Monoxide 



The phosphorus distills as a vapor through a pipe into a 

trough of water, where it collects 





Fig. 62. — Apparatus for the manu- 
facture of phosphorus by the old 
method (final stage). 



Fig. 63. — Electric furnace for 
the manufacture of phos- 
phorus. The raw materials in- 
troduced at A are fed in by 
the screw B, the phosphorus 
vapor escapes at C, and the 
slag is drawn off at D. The 
electrodes are E, E, 



as a heavy liquid. (2) By a new process phosphorus is 
manufactured in an electric furnace (Fig. 63). A mixture; 



PHOSPHORUS 393 

of phosphate, carbon, and sand is fed continuously into a 
furnace provided with an outlet pipe near the upper part 
through which the phosphorus vapor passes into a condenser. 
The residue is drawn off as a slag at the bottom of the fur- 
nace. The process is an electrothermal one, the essential 
equation for the chemical change being — 

2 Ca*(P0 4 )2 + 6 Si0 2 + 10 C = P 4 + 10 CO + 6 CaSi0 3 

Calcium Sand Carbon Phosphorus Carbon Calcium 

Phosphate Monoxide Silicate 

Each method gives a black product, which is purified by 
redistillation in an iron retort, or by oxidation under water 
with sulphuric acid and potassium dichromate; finally it is 
pressed through canvas bags and molded into sticks. 

Properties. — Phosphorus has two allotropic modifica- 
tions, — yellow (waxy) or ordinary and red. Ordinary 
phosphorus when freshly prepared is a yellow, translucent 
solid, but the color deepens by exposure to light. At ordi- 
nary temperatures it is like wax, but at low temperatures 
it is brittle. Under water it melts at 44° C. Exposed to 
the air, it immediately gives off white fumes, and at 34° C. 
takes fire and burns with a brilliant flame, the main product 
being phosphorus pentoxide (P2O5). It is luminous in 
moist air, as may be easily seen by rubbing a phosphorus- 
tipped match in a dark room. This property gave the ele- 
ment its name (from the Greek word phosphoros, light 
bringer). The ease with which it ignites makes phosphorus 
dangerous to handle. Phosphorus is kept beneath water, 
and should never be handled or cut unless so covered. 
Burns from it are severe and hard to heal. It is very poison- 
ous, and the workmen in factories where phosphorus is used 
are liable to contract a dreadful disease, which rots the bones. 
Phosphorus is nearly insoluble in water, but dissolves in 
carbon disulphide and slightly in sodium hydroxide solution. 



394 INORGANIC CHEMISTRY 

Yellow phosphorus has a faint but characteristic odor, 
which may be easily detected by smelling of a phosphorus- 
tipped match. Red phosphorus is made by heating ordi- 
nary phosphorus to about 250° C. in a closed vessel freed 
from air. Conversely, if red phosphorus is distilled and the 
vapor condensed quickly, the yellow variety is obtained. 
This red modification of phosphorus is a dark red powder, 
though sometimes it is a brittle mass. It is opaque and 
odorless, does not glow or take fire when exposed to the 
air, and does not ignite until heated to 260° C. It is not 
poisonous, and does not dissolve in carbon disulphide. It 
can be handled without danger and need not be kept be- 
neath water. Obviously it is the less active variety of the 
element. The specific gravity of red phosphorus is about 
2.2 and that of the yellow form is about 1.83. 

Certain rat and bug poisons contain yellow phosphorus, 
but most of the phosphorus of commerce is consumed in the 
manufacture of matches (see below). 

The vapor density of both yellow and red phosphorus up 
to approximately 1500° C. corresponds to the molecular 
weight 128. Since the atomic weight is 31, the molecular 
formula is P 4 . At higher temperatures partial dissociation 
occurs. The formula of dissolved phosphorus is also P 4 . 

Oxides of Phosphorus. — The two important oxides are 
phosphorous or the trioxide (P 2 3 or P 4 6 ) and phosphoric 
or the pentoxide (P 2 5 or P 4 Oi ). Phosphorous oxide is a 
white solid formed by the slow oxidation of phosphorus or 
by burning phosphorus in a limited supply of air. It has 
the odor of phosphorus and is poisonous. Warmed in the 
air, it changes into the pentoxide. It unites with water to 
form phosphorous acid, thus : ; — 

PA + 3H 2 = 2H 8 PO s 

Phosphorous Phosphorous 

Oxide Acid 



PHOSPHORUS 395 

Phosphoric oxide is a white, snowlike solid formed by burn- 
ing phosphorus in an abundant supply of air. It is very de- 
liquescent and is often used in the laboratory to dry gases. 
It combines vigorously with cold water, forming metaphos- 
phoric acid, thus : — 

P 2 5 + H 2 = 2HP0 3 

Phosphoric Metaphosphoric 

Oxide Acid 

Acids and Salts of Phosphorus. — There are three phos- 
phoric acids, — orthophosphoric (H3PO4) , metaphosphoric 
(HPO3), and pyrophosphoric (H 4 P 2 7 ). Phosphorous acid 
(H3PO3) and hypophosphorous (H 3 P0 2 ) are less important 
compounds. 

Orthophosphoric Acid is a by-product in the manufacture of 
phosphorus from bone ash (see above) ; it can be prepared 
by oxidizing red phosphorus with dilute nitric acid, or by 
dissolving phosphorus pentoxide in hot water, thus : — 



p 2 o 5 + 


3H 2 


= 2H 3 P0 4 


Phosphorus 




Orthophosphoric 


Pentoxide 




Acid 



It is a white deliquescent solid which is veiy soluble in water. 

Metaphosphoric Acid is formed by heating orthophosphoric 
acid to a high temperature, thus : — 

H 3 P0 4 = HPO3 + H 2 

Orthophosphoric Metaphosphoric 

Acid Acid 

It can be formed by dissolving the pentoxide in cold water, 
thus : — 

P 2 5 + H 2 = 2 HPO3 

At ordinary temperature it is a glassy solid, and is called 
glacial phosphoric acid. It dissolves readily in water, and 



396 INORGANIC CHEMISTRY 

the solution changes into orthophosphoric acid — slowly in 
the cold, rapidly when boiled. 

Pyrophosphoric Acid is formed by heating orthophosphoric 
acid or one of its salts to 200°-300° C. ; thus: — 



2H 3 P0 4 = H 4 P 2 7 H 2 

3hos] 
Acid 



Orthophosphoric Pyrophosphoric 
:id Acid 



A sodium salt (HNa 2 P0 4 ) of the ortho-acid is usually used. 
This acid is an amorphous, glassy (but sometimes crystalline) 
solid, readily soluble in water. 

The acids of phosphorus just discussed form salts called 
phosphates, many of which are found as minerals, especially 
phosphates of calcium, aluminium, and magnesium. The 
bones of animals and ashes of plants contain calcium and 
magnesium phosphates. Orthophosphoric acid is tribasic, 
and its salts, which are usually called simply phosphates, 
are numerous. They are known as primary, secondary, 
and tertiary phosphates, according as one, two, or three 
atoms of hydrogen are replaced. The most important is the 
normal calcium salt Ca 3 (P0 4 ) 2 , which has already been de- 
scribed. Hydrogen disodium phosphate (HNa 2 P0 4 ) is the 
commercial sodium phosphate; it is a secondary phosphate. 
This salt and hydrogen sodium ammonium phosphate, or 
microcosmic salt (HNa(NH 4 )P0 4 ), are used in chemical 
analysis. The "acid phosphate " sold as a beverage is a 
solution of one or more acid calcium phosphates (HCaP0 4 
and H 4 Ca(P0 4 ) 2 ). Metaphosphates are formed by heating 
primary (or mono-) sodium phosphates, thus : — 

H 2 NaP0 4 = NaP0 3 + H 2 

Primary Sodium Sodium Meta- 
Phosphate phosphate 

Microcosmic salt when fused also forms a metaphosphate 
(NaP0 3 ) owing to the loss of water and ammonia (NH 3 ). The 



PHOSPHORUS 397 

glassy residue is called a phosphorus bead and like the borax 
bead assumes different colors when heated with metallic 
oxides. Pyrophosphates (of which only two classes are 
known) are formed by heating secondary (or di-) phosphates, 
thus : — 

2HNa 2 P0 4 = Na 4 P 2 7 + H 2 

Disodium Sodium Pyro- 

phosphate phosphate 

Hypophosphites are the salts of hypophosphorous acid 
and are produced by treating phosphorus with an alkali. 
They are often used as medicines. 

Orthophosphoric acid dissociates mainly into the ions 
H + and H 2 P0 4 ". The disodium phosphate (HNa 2 P0 4 ) 
dissociates into the ions 2 Na + and HP0 4 ~" "\ Its solution 
is slightly alkaline because hydroxyl ions are left when the 
hydrogen ion of the slightly dissociated water forms the 
H 2 P0 4 -ion with the HP0 4 -ion, the latter ion having only a 
very slight, if any, tendency to dissociate; the simple ionic 
equation for the hydrolysis is — 

H + + OH- + HPOr - -> oh- + H 2 POr 

Tests for Phosphoric Acids and Phosphates. — Phosphates 
can be distinguished by silver nitrate. Orthophosphates 
give yellow silver phosphate (Ag 3 P0 4 ), metaphosphates give 
white silver metaphosphate (AgP0 3 ), pyrophosphates give 
white silver pyrophosphate (Ag 4 P 2 7 ) ; all dissolve in am- 
monium hydroxide. Metaphosphoric acid coagulates a 
solution of albumin {e.g. white of egg), but orthophosphoric 
and pyrophosphoric acids do not. Orthophosphoric acid 
and its salts precipitate yellow ammonium phosphomolyb- 
date from an excess of a solution of ammonium molybdate. 

Other Compounds of Phosphorus. — Phosphine (PH 3 ) is 
analogous to ammonia (NH 3 ), though it is not alkaline. It 



398 



INORGANIC CHEMISTRY 



is made by heating sodium (or potassium) hydroxide with 
phosphorus. It is poisonous, and has a disagreeable odor; 
it usually takes fire spontaneously in the air owing to the 
presence of an inflammable compound of phosphorus and 
hydrogen. Phosphine itself does not take fire spontaneously. 
It combines with other substances, forming phosphonium 
compounds, which are analogous to ammonium compounds, 
e.g.: — 

PH 3 + HI = PHJ 



Phosphine 



Hydriodic 
Acid 



Phosphonium 
Iodide 



Phosphorus trichloride (PC1 8 ) is a disagreeable smelling liquid, 
made by the combustion of dry chlorine and phosphorus; 
and phosphorus pentachloride (PC1 5 ) is a greenish solid made 
by passing chlorine into a vessel containing the trichloride. 
Both trichloride and pentachloride interact readily with 
water, forming phosphorus compounds and hj^drochloric 
acid, thus: — 



PC1 3 + 


3H 2 


= H3PO3 


+ 3HC1 


Phosphorus 
Trichloride 


Water 


Phosphorous 
Acid 


Hydrochloric 
Acid 


PCI5 + 


H 2 


= POCl 3 


+ 2HC1 


Phosphorus 
Pentachloride 


Water 


Phosphorus 
Oxychloride 


Hydrochloric 
Acid 


PC1 5 + 


4H 2 


"= H 3 P0 4 


+ 5HC1 


Phosphorus 
Pentachloride 


Water 


Phosphoric 
Acid 


Hydrochloric 
Acid 



If either chloride is exposed to moist air, white fumes are 
formed owing to the liberation of hydrogen chloride (HC1). 

When phosphorus pentachloride is heated, it sublimes 
without melting, and under special conditions of temper- 
ature and pressure it becomes a vapor. The molecular 
weight determined from the density of this vapor at about 
300° C. is only about half the calculated value, i.e. 104 in- 
stead of 208.5. Examination of this vapor shows that it is 



PHOSPHORUS 399 

not phosphorus pentachloride, but almost entirely a mixture 
of phosphorus trichloride vapor and chlorine gas. The color 
is greenish and these two components can be separated by 
diffusion. These facts mean that phosphorus pentachloride 
when heated dissociates into phosphorus trichloride and 
chlorine. The reaction is reversible and the equation may 
be written thus : — 

PC1 5 ^ PC1 3 + Cl 2 

At 300° C. equilibrium is maintained by about 3 per cent of 
phosphorus pentachloride and 97 per cent of the trichloride 
and chlorine. This equilibrium may be displaced and the 
reaction sent in the reverse direction by increasing the con- 
centration of one of the reaction products, i.e. by adding 
phosphorus trichloride or chlorine to the tube in which the 
equilibrium prevails. As a matter of fact, when phosphorus 
pentachloride is vaporized in a vessel containing an excess of 
the trichloride vapor, the dissociation of the pentachloride is 
reduced to a minimum, for the vapor density of the penta- 
chloride then corresponds to a molecular weight of 209, the 
calculated value being 208.5. (See Chemical Equilibrium, 
Chapter XI.) 

Matches. — Phosphorus until 1913 was chiefly used in the manu- 
facture of matches. Now a prohibitive tax compels the use of a 
substitute, which is usually a phosphorus sulphide (P 4 S 3 ). In 
making common matches one end of the match stick is first dipped 
into melted sulphur or paraffin and then into the " phosphorus mix- 
ture." The latter consists of different proportions of phosphorus 
sulphide, manganese dioxide (or other oxidizing substance), glue, 
and a little coloring matter. By rubbing them on a rough surface 
enough heat is generated to cause the phosphorus to unite with the 
oxygen of the oxidizing agent, and the heat thereby produced sets 
fire to the sulphur or paraffin, and this in turn kindles the wood. 
In safety matches the head is usually a colored mixture of anti- 
mony sulphide, potassium chlorate, and glue, while the rubbing 
surface is a mixture of red phosphorus, glue, and powdered glass. 



400 INORGANIC CHEMISTRY 

The law imposing the tax on phosphorus matches (two cents 
per hundred matches) was passed mainly to protect workmen from 
the disease caused by breathing fumes of phosphorus. 

Relation of Phosphorus to Life. — Phosphorus is essential 
to the growth of plants and animals. Plants take phosphates 
from the soil and store up the phosphorus compounds, espe- 
cially in the fruit and seeds. Animals eat this vegetable 
matter, assimilate the phosphorus compounds, and deposit 
them in the bones, brain, and nerve tissue. Bones contain 
about 80 per cent of calcium phosphate (Ca 3 (P0 4 ) 2 ). Part 
of the combined phosphorus consumed by animals is re- 
jected by them, and often finds its way back into the soil. 

The constant removal of phosphates by plants would soon 
exhaust the soil. Hence phosphorus is restored to the soil 
in the form of natural or artificial fertilizers. Natural fer- 
tilizers are (1) stable refuse, which always contains some 
of the phosphates from the food originally fed to the ani- 
mals; (2) guano, which is the dried phosphatic and nitrog- 
enous excrement of the sea birds that once lived in vast 
numbers in Peru and Chile ; and (3) phosphate slag, which 
is a phosphorus by-product obtained in manufacturing steel. 
Artificial fertilizers are made from phosphate rock. This 
occurs in large beds in South Carolina, Tennessee, and 
Florida, which yield about a million tons a year. It consists 
of the hardened remains of land and marine animals, and is 
mainly tricalcium phosphate (Ca 3 (P0 4 ) 2 ). It is insoluble 
in water, and must be changed into the soluble primary cal- 
cium salt (H 4 Ca(P0 4 ) 2 ) before it can be easily taken up by 
plants. This soluble salt is called " superphosphate of lime." 
When phosphate rock is treated with sulphuric acid, the 
changes involved may be written thus : — 

Ca 3 (P0 4 ) 2 + 2 H 2 S0 4 = H 4 Ca(P0 4 ) 2 + 2CaS0 4 

Tricalcium " Superphosphate Calcium 

Phosphate of Lime " Sulphate 



ARSENIC 401 

Ca 3 (P0 4 ) 3 + 3H,SO« = 2H 3 P0 4 + 3CaS0 4 

Phosphoric 
Acid 

Ca 3 (P0 4 ) 2 + HaSO* = H 2 Ca 2 (P0 4 ) 2 + CaS0 4 

Dicalcium 
Phosphate 

The aim is to convert the crude phosphate rock into " super- 
phosphate/' but the three reactions usually occur. The 
product is ground, dried, and packed in bags for the market. 
On standing, it may undergo "reversion," i.e. the "super- 
phosphate" and the phosphoric acid may form insoluble 
phosphates, thus making the fertilizer less valuable. Some- 
times "superphosphate" is mixed with compounds of nitro- 
gen and of potassium to produce a complete fertilizer. 

Arsenic 

Occurrence. — Arsenic is found free in nature, but it 
usually occurs combined with sulphur, a metal, or both. 
The common arsenic ores are realgar (As 2 S 2 ), orpiment 
(AS2S3), arsenic pyrites or mispickel (FeSAs). Arsenic 
trioxide or arsenolite (As 2 3 ) is also found. Small quantities 
of arsenic occur in many ores. 

Preparation and Properties. — Arsenic is prepared in the 
laboratory by reducing arsenious oxide with charcoal in a 
glass tube; the arsenic is volatile and is deposited as a dark 
ring on the colder part of the tube. The change is repre- 
sented thus : — 

2As 2 3 + 6C = As 4 + 6CO 

Arsenious Carbon Arsenic Carbon 

Oxide Monoxide 

On a large scale arsenic is extracted from its ores by the 
method just indicated or by roasting arsenic pyrites (FeSAs) 
in the absence of oxygen. 



% 



402 INORGANIC CHEMISTRY 

Arsenic is a brittle, steel-gray solid. A freshly broken piece 
has a metallic luster, which disappears slowly in a moist 
atmosphere. It tends to crystallize. The specific gravity 
varies from 5.62 to 5.96. Heated in the air, it volatilizes 
without melting, and the vapor has an odor like garlic. At 
about 180° C. it burns in the air with a bluish flame, forming 
white arsenious oxide (As 2 3 ). 

The vapor density of arsenic at about 650° C. corresponds 
to the molecular weight 300. Since the atomic weight is 75, 
the molecular formula of arsenic vapor is As 4 at this tempera- 
ture. At about 1700° C. the formula is As 2 . 

Metallic arsenic has few uses, the main one being to harden 
the lead which is made into shot. 

Arsenious Oxide or Arsenic Trioxide, As 2 3 or As 4 Og, is the 

most important compound of arsenic, and is often called 
simply '' arsenic" or " white arsenic." Small quantities are 
found free in nature. The commercial substance is obtained 
as a by-product in the roasting of ores containing arsenic. 
There are two common varieties, a white, granular powder 
and an amorphous, glasslike solid. It is an odorless white 
solid with a faint metallic taste; it dissolves only slightly 
in cold water, but is transformed readily by hot hydro- 
chloric acid into soluble arsenic trichloride (AsCl 3 ). Arsenic 
trioxide is a violent poison. The antidote is fresh ferric 
hydroxide, which may be quickly made by adding ammonium 
hydroxide to a ferric salt, e.g. ferric chloride; the efficiency 
of the antidote depends upon the fact that the ferric hy- 
droxide (formed by the interaction of ammonium hydrox- 
ide and ferric chloride) produces an insoluble compound 
with the arsenic compound. Small doses of arsenic trioxide 
(2 to 3 grains) are usually fatal, but by habitual use the sys- 
tem appropriates larger doses without ill effects. Workmen 
in arsenic factories often accidentally swallow with impunity 



ARSENIC 403 

quantities which would ordinarily prove fatal. It is used 
to a limited extent in making pigments, in manufacturing 
glass and arsenic compounds, in calico printing, in preserv- 
ing skins, and in preparing certain insect and vermin poisons. 
At ordinary pressures arsenic trioxide sublimes without 
melting, and the commercial substance is purified by subliming 
the impure arsenic dust taken from the flues or chambers 
connected with ore furnaces. The vapor density at about 
800° C. requires the formula As 4 6 , but at a very high tem- 
perature the formula is As 2 3 . The formula is also As 4 6 
according to the boiling-point method. 

Other Arsenic Compounds. — The native mineral orpi- 
ment (As 2 S 3 ) is used in making a yellow paint, and realgar 
(AS2S2) a red paint. Arsenic forms acids analogous to the 
acids of phosphorus, though they are less important. Or- 
thoarsenic acid (H 3 As0 4 ) is a white deliquescent solid pre- 
pared by the interaction of concentrated nitric acid and 
arsenic or arsenious oxide. Arsenious acid (H 3 As0 3 ) is known 
only in the solution of its corresponding anhydride (As 2 3 ), 
resembling carbonic acid (H 2 C0 3 ) in this respect. Several 
salts of the acids of arsenic are of interest. Sodium arsenate 
(HNa 2 As0 4 ) and sodium arsenite (Na 3 As0 3 ) are occasionally 
used in dyeing. Scheele's green is chiefly copper arsenite 
(HCuAs0 3 ), and was formerly used to make a cheap green 
paint and to color wall paper. Lead arsenate (Pb 3 (As0 4 ) 2 ) 
and Paris green (a complex compound given the formula 
Cu 3 (As0 3 ) 2 . Cu(C 2 H 3 2 ) 2 ) are effective insecticides and are 
used to exterminate potato bugs and other insect pests. 

Arsenic trisulphide, As 2 S 3 , and pentasulphide, As 2 S 5 , are 
obtained as yellow precipitates by passing hydrogen sul- 
phide gas into acid solutions of arsenious and arsenic com- 
pounds respectively. The formation of this yellow sulphide 



404 INORGANIC CHEMISTRY 

is one test for arsenic. These sulphides form soluble sulpho 
salts with alkaline sulphides, e.g. : — 

As 2 S 3 + 3(NH 4 ) 2 S = 2(NH 4 ) 3 AsS 3 

Arsenious Ammonium Ammonium 

Sulphide Sulphide Sulpharsenite 

Sulpharsenates are formed by ammonium polysulphide, e.g.: 

As 2 S 3 + 3 (NH 4 ) 2 S + S 2 = 2 (NH 4 ) 3 AsS 4 

Ammonium 
Sulpharsenate 

These sulpho-salts are decomposed by hydrochloric acid into 
hydrogen sulphide gas and the yellow sulphides of arsenic. 
If hydrogen sulphide is added to a neutral or basic solution 
of an arsenious compound, the trisulphide is formed as a 
colloid, which is coagulated by adding certain acids or salts. 

Marsh's Test for Arsenic. — Arsenic can be easily detected 
by a simple method, called Marsh's test. An apparatus for 
generating hydrogen is provided with a hard glass horizontal 
delivery tube, narrowed in places and drawn to a point. 
Pure zinc, pure dilute sulphuric acid, and the arsenic solu- 
tion are put in the generator. Hydrogen and gaseous 
hydrogen arsenide (or arsine (AsH 3 )) are formed. If this 
mixture is lighted at the end of the delivery tube, it burns 
with a bluish flame, and metallic arsenic is deposited as a 
black metallic coating on cold porcelain held in the flame; 
or if the tube is heated in front of a narrow place, arsenic is 
deposited at this point. This deposit dissolves in sodium 
hypochlorite solution, but a deposit of antimony, similarly 
produced, does not dissolve. By this delicate test the merest 
trace of arsenic is readily and positively detected. 

Antimony 

Occurrence. — Small quantities of free antimony are 
found. The most common ore is stibnite (Sb 2 S 3 ), which 
occurs in China, Japan, Austria-Hungary, France, Algeria, 



ANTIMONY 405 

Italy, Mexico, and Turkey. Deposits in California and 
Nevada are also utilized. 

Stibnite was found in the fifteenth century. The Latin name of 
antimony is stibium, from stibnite, which gives the symbol of the ele- 
ment, Sb. 

Preparation and Properties. — Antimony is prepared in- 
dustrially by two methods. In one the sulphide is roasted, 
and the oxide thus formed is reduced with charcoal. Equa- 
tions representing the main changes are — 

Sb 2 S 3 ' + 5 2 = Sb 2 4 + 3S0 2 

Antimony Oxygen Antimony Sulphur 
Sulphide Oxide Dioxide 

Sb 2 4 + 4C = 2Sb + 4CO 

The other method consists in heating the sulphide with iron, 
the equation for the chemical change being — 

Sb 2 S 3 + 3Fe = 2Sb + 3FeS 

Antimony Iron Antimony Iron 

Sulphide Sulphide 

Antimony is a silver-white, crystalline, brittle solid. Its 
specific gravity is 6.7. At ordinary temperatures antimony 
does not tarnish in the air, but when heated, it burns with a 
bluish flame, forming the white, powdery antimony trioxide 
(Sb 2 3 ). The melting point is about 630° C. Powdered 
antimony burns brilliantly when added to chlorine, bromine, 
or iodine, owing to the vigorous and rapid combination with 
these elements. Nitric acid oxidizes it to Sb 2 3 or to anti- 
monic acid (H 3 Sb0 4 ), and aqua regia converts it into anti- 
mony trichloride (SbCl 3 ). It is one constituent of type 
metal (see Alloys of Lead) and other alloys. 

Compounds of Antimony. — Antimony forms stibine 
(SbH 3 ), which is analogous to ammonia (NH 3 ) and arsine 
(AsH 3 ), pyro- and meta- acids, the oxides, Sb 2 3 , Sb 2 4 , Sb 2 5 , 



406 INORGANIC CHEMISTRY 

and halogen compounds. It also forms complex compounds 
in which antimony acts as a metal. Tartar emetic, or potas- 
sium antimonyl tartrate (KSbO . C 4 H 4 6 ), is a white solid, 
soluble in water; it is used as a medicine and as a mordant 
in dyeing cotton. (The group SbO is called antimonyl.) 
Antimony trisulphide (Sb 2 S 3 ) as prepared in the laboratory is 
an orange red solid; it is formed by passing hydrogen sul- 
phide gas into a solution of antimony — the test for anti- 
mony. The native sulphide, or stibnite, is a lustrous, blue- 
gray mineral, often beautifully crystallized. Antimony 
chloride (SbCl 3 ) is formed by the action of chlorine upon the 
metal or by boiling the metal in aqua regia; it hydrolyzes 
readily, forming a white solid — antimony oxychloride 
(SbOCl). The formation of antimony oxychloride is some- 
times used as a test for antimony, but the more common 
test is the formation of the orange red sulphide (Sb 2 S 3 ). 
The sulphides of antimony (Sb 2 S 3 and Sb 2 S 5 ) form salts with 
alkaline sulphides, which behave like the corresponding 
arsenic compounds. 

Bismuth 

Occurrence. — Bismuth is usually found in the native 
state, though it is not abundant nor widely distributed. 
The oxide (Bi 2 3 ) or bismite, and the sulphide (Bi 2 S 3 ) or bis- 
muthinite, are the common ores. The world's supply comes 
from Saxony. 

Preparation and Properties. — Bismuth is prepared from 
the native metal by melting it on an inclined plate and allow- 
ing it to drain away from the solid impurities. Sometimes 
the sulphide is roasted and the resulting oxide is reduced 
with charcoal, as in the case of antimony. 

Bismuth is a gray-white metal with a red tinge. Like 
antimony, it is hard and brittle. It does not tarnish in dry 



BISMUTH 407 

air, but it grows dull in moist air; and when heated in air it 
burns with a bluish flame, forming the yellowish oxide (Bi 2 3 ) 
Its specific gravity is about 9.9. Hydrochloric acid does 
not readily attack it; but nitric acid converts it into a nitrate, 
hot sulphuric acid into a sulphate, and aqua regia into a 
chloride. 

Bismuth melts at about 270° C. But a mixture of bismuth, 
lead, and tin melts at a low temperature. For example, 
Newton's metal melts at 95° C. and Rose's metal at 100° C; 
while Wood's metal, which contains cadmium, melts at only 
66°-71° C. These metallic mixtures are called fusible metals 
or alloys. They are used in making casts of wood cuts; but 
more often as safety plugs in steam boilers to prevent ex- 
plosions, as a fuse in electrical apparatus to prevent a short 
circuit, and as a link to hold in place fireproof doors and the 
valves in the automatic sprinkling apparatus now installed 
in large buildings. 

Compounds of Bismuth. — Bismuth forms no compounds 
with hydrogen. There are four oxides, but only two are 
well known. Bismuth trioxide (Bi 2 3 ) is a pale yellow 
powder, and the pentoxide (Bi 2 9 5 ) is a brown powder. Bis- 
muth trioxide is used to fix the gilding on porcelain. The 
trichloride (BiCl 3 ) is formed by treating bismuth with aqua 
regia; it hydrolyzes with an excess of water, forming basic 
bismuth chloride (Bi(OH) 2 Cl), which by loss of water be- 
comes bismuth oxy chloride (BiOCl). The formation of the 
white insoluble oxychloride is a test for bismuth. Bismuth, 
being a metal, forms hydroxides (Bi(OH) 3 and BiO . OH). 
Normal bismuth nitrate (Bi(N0 3 )3) when treated with hot 
water forms basic bismuth nitrate (Bi(OH) 2 N0 3 or BiON0 3 ). 
The latter, often called subnitrate of bismuth, is a white 
powder, and is used as a medicine for dyspepsia and as a 
cosmetic. 



408 INORGANIC CHEMISTRY 

The Nitrogen Family. — This family illustrates typically 
the relation of atomic weight to properties, for these elements 
display a gradual change from non-metal to metal as we pass 
from nitrogen (atomic weight 14.01) through the periodic 
members to bismuth (atomic weight 208.0). Nitrogen and 
phosphorus are distinctly non-metallic, arsenic is both non- 
metallic and metallic, antimony is increasingly metallic, 
while bismuth is a typical metal. They form analogous 
compounds, for example : — 



NH 3 


N 2 3 


N 2 O s 




PH S 


p 2 o 3 


PA 


PC1 3 


AsH 3 


As 2 3 


As 2 O s 


AsCl 3 


SbH 3 


Sb 2 3 N 


Sb 2 5 


SbCl 3 





Bi 2 3 


Bi 2 O s 


BiCl 3 



Problems and Exercises 

1. Calculate the percentage composition of (a) sodium phosphate 
(Na3P0 4 ), (b) dihydrogen sodium phosphate (H 2 NaP0 4 ), (c) disodium 
phosphate (HNa 2 P0 4 ), (d) microcosmic salt (HNaNH 4 P0 4 ). 

2. How much phosphorus is needed to remove the oxygen from a 
liter of air? (Assume (1) 2 P + 5 O = P2O5, and (2) air is 20 per cent 
oxygen.) 

3. How much phosphorus is there in a ton (2000 lb.) of bone ash 
if 70 per cent of the sample is calcium phosphate (Ca3(P04)2)? 

4. If a skeleton weighs 25 lb. and contains 60 per cent calcium 
phosphate, how much phosphorus does it contain? 

5. What is the weight of a cylindrical stick of ordinary phosphorus 
10 cm. long and 15 mm. in diameter? (Suggestion. — What is the 
specific gravity of phosphorus?) 

6. Calculate the percentage composition of (a) orpiment (AS2S3), 
(b) realgar (AS2S2), (c) white arsenic (AS2O3). 

7. What is the weight of a piece of antimony 25 cm. long, 15 cm. 
wide, and 2 mm. thick? 

8. What volume of chlorine and of phosphorus trichloride is formed 
by the complete dissociation of 20 1. of phosphorus pentachloride ? 

9. If 18.5854 gm. of phosphorus yield 42.584 gm. of phosphorus 



BISMUTH 409 

pentoxide, what is the atomic weight of phosphorus? (At. wt. of 
oxygen =16.) 

10. If 2.99091 gm. of antimony combine with 1.19495 gm. of sul- 
phur to form antimony sulphide (Sb 2 S3), what is the atomic weight 
of antimony if the atomic weight of sulphur is 32.07? 

11. Suppose 29.5305 gm. of bismuth trioxide (Bi 2 3 ) yield 3.044 gm. 
of oxygen. If 16 is the atomic weight of oxygen, what is the atomic 
weight of bismuth? 

12. Write the formulas of the following compounds by applying 
the principle of valence (see Chapter XIV): (a) Arsenious chloride, 
arsenic chloride, hydrogen arsenide, arsenic iodide; (b) hydrogen 
antimonide, antimonic chloride, antimonious sulphide; (c) bismuth 
hydroxide, bismuth nitrate, bismuth sulphate; (d) calcium phos- 
phide, acid calcium phosphate, phosphorous bromide, phosphoric 
iodide, magnesium phosphate (ortho), ammonium magnesium phos- 
phate. 

13. Calculate the weight of phosphorus theoretically obtain- 
able from (a) 2 metric tons of phosphorite (70 per cent Ca 3 (P0 4 )2), 
(b) 3 metric tons of apatite (90 per cent pure), and (c) 4 metric tons 
of bone ash (82 per cent Ca 3 (P0 4 )2). 

14. How much (a) metaphosphoric acid and (6) pyrophosphoric 
acid can be made from 200 gm. of orthophosphoric acid (90 per 
cent pure) ? 

15. Calculate the per cent of (a) antimony in an ore containing 
37 per cent of stibnite, and (b) of bismuth in an ore containing 45 
per cent of bismite. 

16. Calculate the simplest formulas corresponding to (a) P 
= 25.833, = 53.333, H= 1.666, Na= 19.166; (b) P = 21.83, 
= 45.07, H=.70, Na= 32.39; (c) As = 61, S = 39; and (d) As 
= 70.1, S = 29.9. 

17. Write the volumetric equations for (a) formation of phos- 
phorus pentoxide from phosphorus vapor and oxygen, and (b) prepa- 
ration of phosphorus trichloride, and (c) preparation of phosphorus 
pentachloride. 

18. Phosphine (PH 3 ) burns to phosphorus pentoxide and water. 
What volume of oxygen is needed for 800 cc. of phosphine? 



CHAPTER XXIV 
Metals and Metallurgy 

Introduction. — The elements studied thus far are chiefly 
non-metals. Metals, however, have been mentioned, and 
many of their properties have been discussed. In the present 
chapter we shall review these properties and prepare the way 
for a fuller treatment of the metals. 

Metals and Non-metals. — Many years ago the chemical 
elements were divided into two classes, called metals and 
non-metals. The division was based largely on the conspicu- 
ous physical properties of the elements. The opaque, lus- 
trous, more or less heavy, hard, ductile, malleable, tenacious 
solids were called metals. All gases and certain solids, such 
as carbon, sulphur, phosphorus, and iodine, were called 
non-metals. Their chemical properties do not permit such a 
sharp dividing line, however, to be drawn between metals 
and non-metals. Some elements have pronounced properties, 
like the non-metal sulphur and the metal iron; these are typ- 
ical. A few elements have variable properties; sometimes 
they act as metals and sometimes as non-metals. Antimony 
and arsenic belong to this border-line class; they are occa- 
sionally called the metalloids. The classification into metals 
and non-metals is not accurate from a strictly chemical stand- 
point, but it is serviceable. The use in common life of the 
words metallic and metal seldom leads to confusion. 

Although lists of the metals and non-metals have already 
been given, a repetition on a slightly different basis is not 
inappropriate in this chapter. The following is a 

410 



METALS AND METALLURGY 



411 



O 






- 



Iron 

Cobalt 

Nickel 

Platinum 






<d 

GO 

0) 

bD 


Fluorine 
Chlorine 
Bromine 
Iodine 


CD 
CO 

CD 

eel 
bD 

5 


O 

5 


Oxygen 
Sulphur 


3 

o 

s-l 

5 


>> 


Nitrogen 

Phosphorus 

Arsenic 


O £ 

1 1 

< < 


.5 © 


Carbon 
Silicon 


3 


s 
5| 


o 
u 
o 
pq 


B 
< 


Calcium 

Strontium 

Barium 

Magnesium 
Zinc 

Cadmium 
Mercury 






Sodium 
Potassium 

Copper 

Silver 

Gold 


CD 

bD 

p 




CD 

a 


© 
o 

ft 


XJl 

"o 

CD 



412 INORGANIC CHEMISTRY 

Physical Properties of Metals. — The physical properties 
of most metals are familiar. The properties of individual 
metals vary somewhat, depending upon the method of manu- 
facture and the temperature. Until recently some metals 
were known only as powders, but the electrothermal and 
aluminothermal processes now enable us to prepare most 
of them in coherent masses. All metals have a metallic 
luster, i.e. the marked property of reflecting light from 
their polished or untarnished surfaces. All are opaque 
except in very thin films. The color of many is white, 
though the tint varies. Thus, in a compact state silver, 
sodium, aluminium, mercury, magnesium, iron, and tin 
are nearly pure white, while bismuth is reddish white and 
potassium is bluish. Copper is the only red metal and gold 
the only yellow one which is an element. When powdered, 
several metals are dark, some even black. Most metals 
are malleable and ductile, i.e. they can be hammered or rolled 
into sheets and drawn into wire. Gold, copper, silver, iron, 
platinum, and aluminium possess both these properties to 
a marked degree; while lead, tin, and zinc are very malle- 
able though not especially ductile. Antimony and bismuth 
are brittle. The hardness of metals varies. At the ordinary 
temperature mercury is a liquid, sodium and lead can be 
cut easily with a knife, and so on through the list up to 
iridium, which is as hard as steel. In specific gravity, 
which was once thought to be very high, the metals range 
between lithium, which has the specific gravity .534, and 
osmium, which has the specific gravity 22.48. Sodium 
and potassium are lighter than water, while magnesium has 
the specific gravity 1.75, and aluminium 2.58. Metals 
conduct heat and electricity. They also vary in this prop- 
erty. Silver, copper, and aluminium are the best con- 
ductors, and have therefore many practical applications. 
Bismuth is the poorest conductor. 



METALS AND METALLURGY 413 

Chemical Properties of Metals. — Bases are formed when 
oxides of metals dissolve in water. On the other hand 
acids result from dissolving oxides of non-metals. Thus : — 

CaO + H 2 = Ca(OH) 2 

Calcium Calcium 

Oxide Hydroxide 

S0 3 + H 2 = H^SO, 



Sulphur Sulphuric 

Trioxide Acid 



/ 



Metals, therefore, are base-forming elements. When com- 
pounds of metals are dissolved in water, the metal becomes 
the positive ion or cation, and the solutions have properties 
characteristic of the metal in the ionic state. Thus, solu- 
tions of copper nitrate and copper sulphate respond to 
the same test for copper because both contain copper ions 
(Cu ++ ). Sometimes the solution contains the metal as 
part of a compound ion, e.g. in solutions of potassium 
ferrocyanide the iron is present as the ion Fe(CN) 6 . 

Occurrence of Metals. — Only a few metals are found 
uncombined or free in the earth's crust, and these are seldom 
pure. Of the six metals known to the ancients (gold, copper, 
silver, tin, iron, and lead) all except tin and lead are found 
free. The metals which occur free in the earth's crust are 
called native, while their compounds are called minerals; 
the term mineral, however, is also applied to certain in- 
organic substances found in the earth's crust {e.g. sulphur, 
graphite, silica). Those minerals from which metals can 
be profitably extracted are called ores; sometimes the 
term ore is also applied to a rock containing native metals, 
e.g. gold ore or copper ore. The most abundant classes of 
ores are oxides, sulphides, carbonates, sulphates, and 
hydroxides. Many ores contain arsenic. Some ores are 
very complex. 



414 INORGANIC CHEMISTRY 

Preparation of Metals. — The series of operations by which 
useful metals are extracted from their ores is called met- 
allurgy. It includes preliminary treatment, smelting, elec- 
trolysis, refining, and other necessary operations. 

The object of the preliminary treatment is to prepare 
the ore for smelting or for a similar operation. The ore 
as it comes from the mine is usually mixed with earthy 
matter or rock called gangue. This impurity is removed 
by mechanical or chemical processes, sometimes by both. 
The mechanical process illustrates one kind of preliminary 
treatment. The ore is first crushed in a stamp mill. This 
is a huge, heavy mortar and pestle. The pestle falls re- 
peatedly upon the ore, which is slowly fed into the mortar 
or die. A current of water (or air) forces the fine particles 
out of the mortar through a sieve. The lighter particles 
of the impurities are washed away, and the metallic grains 
are extracted by another mechanical operation, though 
chemical processes are frequently employed, especially with 
inferior ores. This separation of the valuable part of the 
ore from the gangue and reducing it to a smaller bulk is 
often called ore dressing or concentration. Copper is ex- 
tracted from Lake Superior ores mainly by this method of 
preliminary treatment. 

Gold and silver ores are often treated like copper and then 
extracted from the slime by mercury. The latter operation 
is called amalgamation. The most common method of 
extracting metals from their ores is by smelting. The 
process varies with the kind and composition of the ore. 
Essentially, it consists in heating a mixture of the ore and 
coke (or coal) in a furnace. The ores used must, as a rule, 
be oxides. Sulphides, hydroxides, and carbonates are first 
roasted or calcined to convert them into oxides. The 
essential chemical change in smelting is a reduction of the 
oxide by carbon. The carbon and oxygen unite and pass off 



METALS AND METALLURGY 



415 



as a gas, leaving the molten metal at the bottom of the 
furnace. Limestone or a similar substance is often added 
to the mixture as a flux, i.e. to facilitate the melting and to 
assist in removing the impurities as a glassy substance, 
called slag. The operation is conducted in different kinds 
of furnaces. Iron, for example, is melted in a huge upright 
furnace called a blast furnace (Fig. 80), because a current 
of air is forced through the melted mass to facilitate the 
fusion and chemical changes. In such a furnace the fuel 
and ore are in direct con- 
tact. When this is unde- 
sirable, the reverberatory 
furnace is used (Fig. 64). 
As the figure shows, in 
this furnace the flame is 
reflected or reverberated 
upon the ore under treat- 
ment. In this kind of 
furnace the ore may be oxi- 
dized or reduced without 
coming in contact with the 
fuel. Some ores demand 
special methods, which 
will be described in con- 
nection with the metals. 

Electrolysis is used to extract some metals, especially 
aluminium and magnesium. Gold, silver, lead, and copper, 
are purified by electrolysis. A few metals are extracted by 
a wet process. That is, the ores are dissolved, and the 
metal is then precipitated by adding some substance or by 
electrolysis. Thus, inferior gold ores are dissolved by treat- 
ment with potassium (or sodium) cyanide, and the gold is 
then precipitated by zinc. 

Some metals, hitherto rare, are obtained by reducing their 




Fig. 64. — Reverberatory furnace. The 
fire burns on the grate G and the long 
flame which passes over the bridge E is 
reflected down by the sloping roof upon 
the contents of the furnace. Gases 
escape through /. The charge, which 
rests upon the bed B, does not come 
in contact with the fuel, but is oxidized 
or reduced by the flame. 



416 INORGANIC CHEMISTRY 

oxides with carbon in the electric furnace or by heating the 
powdered oxide with aluminium, e.g. chromium. 

Alloys are mixtures or compounds of two or more metals. 
Some fused metals mix in all proportions, while others seem 
to form definite compounds. The properties of alloys vary 
with the constituents and their proportions. Some alloys, 
especially those of copper and of lead, have many industrial 
uses. Alloys in which mercury is a constituent are called 
amalgams. 

Problems and Exercises (Review) 

1. What is the specific gravity of gold, if a piece weighs 4.676 
gm. in air, and loses 0.244 gm. when weighed in water? (Note. — 
Specific gravity equals the weight in air divided by the loss of 
weight in water.) 

2. The weight of a liter of ether vapor at 100° C. and 760 mm. 
is 2.44 gm. What is the molecular weight of ether? 

3. (a) What is the atomic weight of phosphorus, if the specific 
heat is .189? (6) Of potassium, if the specific heat is .166? (c) Of 
manganese, if the specific heat is .122? 

4. If a liter of neon (at 0° C. and 760 mm.) weighs .902 gm., 
what is the atomic weight of this monatomic gas? 

5. Complete and balance the following equations : (a) Cu ++ + 

+S-- + = CuS+H++N0 3 -; (b) Ba++ + — -- h S0 4 — 

+ = +2H++C1- 

6. Calculate the weight of zinc dissolved by 100 gm. of a solu- 
tion of hydrochloric acid containing 20 per cent by weight of HC1. 

7. What volume of acetylene (standard conditions) will 200 
pounds of calcium carbide yield? 

8. Calculate the formulas corresponding to (a) N = 26.168, 
CI =66.355, H = 7.476; (b) N = 22.222, 0-76.19, H= 1.587; 
(c) N = 16.47, O == 56.47, Na = 27.06. What is the name of each 
compound ? 

9. What is the weight of sulphur in 20 1. of sulphur dioxide 
measured at 20° C. and 780 mm. pressure? 

10. Name the sodium salt of (a) nitrous and (6) nitric acid. 
Name the corresponding salts of K, barium, Ca, silver, Pb, zinc, 
NH 4 , aluminium. 



CHAPTER XXV 

Sodium, Potassium, Lithium, and Ammonium — Spectrum 

Analysis 

Introduction. — Sodium and potassium, together with the 
less common element lithium and the rare elements rubid- 
ium and caesium, form a natural family in Group I of the 
periodic classification, known as alkali metals. These 
elements and their corresponding compounds resemble 
each other closely. 

Compounds of the hypothetical metal ammonium are 
conveniently treated in this chapter because their chemical 
relations are similar. 

Sodium and potassium were discovered by Sir Humphry Davy 
in 1807 by the electrolysis of their hydroxides. Bunsen, by means of 
the spectroscope, discovered lithium in 1855, caesium in 1860, and 
rubidium in 1861. 

Sodium 

Occurrence. — Sodium is not found free. Sodium chloride 
and sodium nitrate are the most abundant compounds. 
Many minerals, rocks, marine plants, and mineral waters 
contain combined sodium. About 2.5 per cent of the earth's 
crust is sodium. 

The symbol of sodium, Na, is from the Latin word natrium, which 
in turn comes from the Greek word natron, an old name of sodium 
carbonate. 

Preparation. — Sodium is manufactured on a large scale 
by the electrolysis of fused sodium hydroxide. This method, 

417 



418 



INORGANIC CHEMISTRY 



£;,,,. ,~ 



z~. 



n 



LD^ 




cc 



£J 



used on a small scale by Davy in 1807 to isolate sodium, 
became practicable only recently and is known as the Castner 

method. Figure 65 is a sketch 
of the apparatus. The body of 
the steel cylinder rests within a 
heated flue. The iron cathode 
(C) passes up through the bottom 
of the cylinder into the fused 
sodium hydroxide. A cylindrical 
collecting pot (P) terminating in 
a wire gauze surrounds the end of 
the cathode. Several carbon bars 
(A, A) dip into the vessel from 
above, and constitute the anode. 
As the electrolysis proceeds, so- 
dium and hydrogen are liberated 
at the cathode and oxygen at 
the anode. The oxygen escapes 
through a pipe (0) without com- 
ing in contact with the sodium, 
while the sodium and hydrogen collect in P. The hydrogen 
escapes through the top of P, while the sodium, which is 
protected from the oxidizing action of the air by the hydro- 
gen, is ladled out at intervals. 



ni 



B 



U 



n>0 



Fig. 65. — Apparatus for the 
manufacture of sodium by 
the electrolysis of sodium 
hydr oxide. 



Properties. — Sodium is a silver-white metal. It can be 
easily cut with a knife and molded with the fingers. It is 
light enough to float upon water, since its specific gravity 
is 0.9712 (at 20° C). Heated in the air, it melts at about 
96° C, and at a higher temperature it burns with a brilliant 
yellow flame, forming sodium peroxide (Na 2 2 ). This 
intense yellow color is characteristic of sodium and is a 
test for the element (free or combined). In moist air the 
bright surface quickly tarnishes, and sodium as usually seen 



SODIUM 419 

has a yellow or gray-brown coating. It is, therefore, kept 
under kerosene or a liquid free from water. 

It decomposes water at ordinary temperatures, liberating 
hydrogen and forming sodium hydroxide, thus: — 

2 Na + 2 H 2 = 2 NaOH + H 2 

Sodium Water Sodium Hydrogen 

Hydroxide 

If held in one place upon water by filter paper, enough heat 
is generated to set fire to the hydrogen, which burns with a 
yellow flame, owing to the presence of volatilized sodium. 
If sodium is melted in chlorine, the two elements combine 
with a brilliant flame, forming sodium chloride. Davy, 
in 1810, proved in this way that common salt is really nothing 
but sodium chloride. It combines with hydrogen and forms 
a white solid called sodium hydride (XaH). If mixed with 
mercury, it forms sodium amalgam, which is sometimes 
used instead of sodium itself. 

A molecule of sodium has been found to be monatomic 
by the vapor density and the freezing-point methods. 

Sodium is used in the laboratory to extract water from 
alcohol and ether and to prepare organic compounds. Large 
quantities are consumed in the manufacture of sodium 
peroxide (Na 2 2 ) and sodium cyanide (XaCX). Its power to 
reduce oxides gives it limited use in preparing certain rare 
metals, e.g. zirconium, tantalum, niobium, and thorium, 
though it is being replaced by aluminium. (See Thermit.) 

Sodium Chloride, NaCl, is the most important compound 
of sodium. It is one of the most abundant substances, 
and is familiar under the name of salt, common salt, or 
table salt. The presence of salt in the ocean, in lakes and 
springs, and in the soil is mentioned in the oldest historical 
records. Sodium chloride constitutes about 77 per cent 



420 INORGANIC CHEMISTRY 

of the salts found in sea water and by far the largest part 
of the salt deposits in the earth's crust. 

Preparation of Salt. — The chief sources of salt are sea water, 
rock salt deposits, and brines. (1) In warm countries, as on the 
shores of the Mediterranean Sea, shallow ponds of sea water near the 
shore are evaporated by exposure to the sun and wind, and the salt 
is collected. In some regions sea water is first concentrated by allow- 
ing it to trickle over heaps of brush and then evaporate to crystalliza- 
tion in shallow pans. In cold countries, as on the shores of the White 
Sea in Russia, sea water is allowed to freeze and the ice is removed. 
The ice contains no salt, so the operation is repeated until the remain- 
ing liquid becomes concentrated enough to evaporate profitably 
over a fire. (2) Deposits of salt are found in many parts of the globe, 
the most important being in England, Austria-Hungary, and Ger- 
many. In these regions and some parts of the United States the salt 
is mined and purified like other minerals. This variety is coarse and 
often impure, and is largely used in curing meat and preserving hides. 
(3) Most of the salt produced in the United States is obtained from 
natural or artificial brines, i.e. from concentrated solutions of salt. 
Artificial brines are made by forcing water into salt deposits. Brines 
are obtained in New York, Michigan, Kansas, Ohio, West Virginia, 
California, Utah, and Louisiana. They are evaporated in vats by 
the sun's heat or by heating in kettles or pans. 

All these methods give a product containing as impurities salts of 
calcium and magnesium, which are largely removed by further special 
treatment. According to the standard established by the United 
States Department of Agriculture, dry table or dairy salt must not 
contain over 1.4 per cent of calcium sulphate, .5 per cent of calcium 
and magnesium chlorides, and .1 per cent of matter insoluble in 
water. The dampness of salt is due to traces of magnesium and 
calcium chlorides. (See Deliquescence.) 

Properties and Uses of Salt. — Salt is rather soluble in 
water, 100 gm. of water dissolving about 36 gm. of salt at 
0° C. and 40 gm. at 100° C. It crystallizes in cubes, which 
often snap open sharply (i.e. decrepitate) when heated, 
owing to the sudden vaporization of the inclosed water. 
This substance is an essential ingredient of the food of man 
and animals. Besides its universal domestic use, enormous 



SODIUM 421 

quantities are consumed in the preparation of many sodium 
and chlorine compounds, especially sodium carbonate, hydro- 
chloric acid, and bleaching powder. 

Sodium Carbonate, Na 2 C0 3 , is next to sodium chloride in 
importance. Small quantities of hydrated sodium car- 
bonates are found in Egypt, Russia, and in California and 
Nevada. Formerly it was obtained from the ashes of marine 
plants, but sodium chloride is now the source. The manu- 
facture of sodium carbonate is one of the most extensive 
chemical industries. Two processes are used, the Leblanc 
and the Solvay. 

The Leblanc Process has three stages. (1) Sodium chloride is 
changed into sodium sulphate by sulphuric acid, the equation for the 
change being — 

2NaCl + H 2 S0 4 = Na 2 S0 4 + 2 HC1 

Sodium Sulphuric Sodium Hydrochloric 
Chloride Acid Sulphate Acid 

This operation is called the salt cake process; the impure product, 
called "salt cake/' contains about 95 per cent of sodium sulphate. 
The hydrochloric acid is a profitable by-product. (See Hydrochloric 
Acid.) (2) and (3) The sodium sulphate is reduced to sodium sul- 
phide by heating the "salt cake" with coal; and the resulting sodium 
sulphide is changed into sodium carbonate by heating with limestone. 
These two chemical changes, which are accomplished by one operation, 
are represented by the following equations : — 

Na 2 S0 4 + 2C = Na 2 S + 2 C0 2 

Sodium Carbon Sodium Carbon 

Sulphate Sulphide Dioxide 

Na 2 S + CaC0 3 = Xa 2 C0 3 + CaS 

Sodium Limestone Sodium Calcium 

Sulphide Carbonate Sulphide 

This operation is called the black ash process. The product is a dark 
brown or gray porous mass, and contains, besides 37 to 45 per cent 
of sodium carbonate, considerable calcium sulphide and other im- 
purities. The sodium carbonate is rapidly separated from the in- 
soluble portions of the " black ash" by extraction with a regulated 



422 INORGANIC CHEMISTRY 

stream of water. The concentrated solution of sodium carbonate thus 
obtained is evaporated to crystallization, and the crude crystals are 
ignited. This product is known as soda ash, and from its solution in 
water are obtained soda crystals or sal soda (Na2CC>3 . IOH2O). 

The Solvay Process, often called the ammonia-soda process, con- 
sists in saturating a cold concentrated solution of sodium chloride first 
with ammonia gas and then with carbon dioxide gas. The equation 
for the complete chemical change is — 



H 2 + NaCl + NH 3 + C0 2 = 


= HNaC0 3 


+ NH4CI 


Water Sodium Ammonia Carbon 
Chloride Dioxide 


Acid Sodium 
Carbonate 


Ammonium 
Chloride 



The acid sodium carbonate is nearly insoluble in the cold ammonium 
chloride solution, and therefore separates. It is changed, by heating, 
into sodium carbonate, thus : — 



2 HNaC0 3 


= Na 2 C0 3 + C0 2 + H 2 


Acid Sodium 
Carbonate 


Sodium Carbon Water 
Carbonate Dioxide 



The liberated carbon dioxide is used again, and from the ammonium 
chloride the ammonia is also recovered and used. 

Properties and Uses of Sodium Carbonate. — Crystallized 
sodium carbonate (Na 2 C0 3 . 10 H 2 0) is often called alkali 
or soda. It loses water in the air, becoming dull at first 
and finally falling to a powder, owing to the fact that its 
vapor pressure is greater than the average vapor pressure 
of the air; this phenomenon, as already stated, is called 
efflorescence. When heated, it melts in its water of crys- 
tallization, and finally becomes the anhydrous salt (soda 
ash, calcined soda, Na 2 C0 3 ). It is readily soluble in water, 
and the solution, which is strongly alkaline, is widely used 
as a cleansing agent, hence the name washing soda. A 
water solution of sodium carbonate is alkaline owing to 
hydrolysis. Sodium carbonate ionizes into 2 Na + and 
C0 3 ~~, but the C0 3 -ions are unstable and form HC0 3 -ions 
with the H-ions from the slightly dissociated water; this 
removal of H-ions finally leaves in the solution sufficient 
OH-ions to produce an alkaline reaction to litmus. 



SODIUM 423 

Besides its general use as a cleansing agent, enormous 
quantities of sodium carbonate are consumed in the glass 
and soap industries and in preparing sodium compounds. 

Sodium Bicarbonate, HNaC0 3 , is formed at one stage of 
the Solvay process (see above). It may also be prepared by 
treating crystallized sodium carbonate with carbon dioxide 
gas. It is a white solid, and is less soluble in water than 
the normal carbonate. When heated or when mixed with 
an acid or an acid salt, sodium bicarbonate gives up carbon 
dioxide. This property early led to its use in cooking, 
and gave the names cooking soda, baking soda, or simply 
soda. 

Sodium bicarbonate is so called because in it the ratio of 
the C0 3 to the Na is twice that in the normal carbonate 
(Xa 2 C0 3 ). Although called acid sodium carbonate, a solu- 
tion of the pure salt is practically neutral, the ions being 
Na+ and HCO3-. 

Sodium bicarbonate is one ingredient of baking powder 
and of various mixtures (except yeast) used to raise bread, 
cake, and other food. The other essential ingredient is 
cream of tartar (acid potassium tartrate), which has a mild 
acid reaction and thus liberates carbon dioxide from the 
bicarbonate (see p. 320). Sour milk, which contains lactic 
acid, is sometimes used in place of cream of tartar. When 
pastry is raised with baking soda and cream of tartar, the 
escaping carbon dioxide puffs up the dough. Hence baking 
soda is often called saleratus — the salt which aerates 
(from the Latin words sal, salt, and aer, air or gas). Ef- 
fervescing powders, such as Seidlitz powder, are mixtures 
of sodium bicarbonate and tartaric acid or one of its acid 
salts. When the ingredients are put into water, carbon 
dioxide is liberated. Sodium bicarbonate is used as a 
medicine to neutralize an acid stomach. For example, the 



424 INORGANIC CHEMISTRY 

" soda mints " sometimes taken for this purpose are mainly 
sodium bicarbonate. 

Sodium Hydroxide or Caustic Soda, NaOH, is a white 
deliquescent, corrosive solid. It absorbs water and carbon 
dioxide rapidly from the air. When exposed to the air, it 
becomes moist at first, then forms a concentrated solution, 
and ultimately solidifies, owing to transformation into so- 
dium carbonate. The deliquescence of sodium hydroxide, as 
already stated, is due to the fact that the solution formed 
from the water deposited on its surface has a lower vapor 
pressure than the vapor pressure of the water vapor in the 
air; hence the solid continues to dissolve in the absorbed 
water until a solution is produced whose vapor pressure 
equals the pressure of the water vapor in the surrounding 
air. It dissolves readily in water with rise of temperature, 
and the solution is strongly alkaline owing to the high degree 
of ionization of the solute. When heated, it melts easily, 
and is often cast into sticks for use in the laboratory. 

Immense quantities are used in making hard soap, paper, 
and dyestuffs ; in bleaching, and in refining kerosene oil. 

Sodium hydroxide is manufactured by two general methods, 
one electrolytic and the other purely chemical. In the elec- 
trolytic method the sodium liberated from sodium chloride 
interacts with water and forms sodium hydroxide. 

The apparatus used in one electrolytic process is shown 
in Figure 66. It consists of a slate box divided into one 
cathode and two anode compartments by partitions extend- 
ing nearly to the bottom ; the compartments are separated 
by a layer of mercury (shown in black). The T-shaped 
anodes (A, A) of graphite and the cathode (C) of iron reach 
nearly to the mercury. The anode compartments contain 
sodium chloride solution, while the cathode compartment 
contains sodium hydroxide solution; sodium chloride solu- 



SODIUM 



425 



tion of the right concentration flows slowly and continuously 
through the anode compartments by means of the pipes E, 
E (and outlets not shown). When the current passes, 
chlorine is evolved at the anodes and escapes through the 




Fig. 66. — Apparatus for the manufacture of sodium hydroxide by the 
electrolysis of sodium chloride. 

pipes D, D ; the sodium is liberated at the intermediate 
cathode of mercury and forms an amalgam with it. By 
carefully rocking the cell on the device X, X, the sodium 
amalgam alone flows beneath the partitions into the cathode 
compartment, where the sodium is liberated at the iron 
cathode ; the sodium at once reacts with the water, forming 
hydrogen and sodium hydroxide. The hydrogen escapes 
through the pipe H, while the sodium hydroxide solution is 
drawn off (through G) and replaced by water (through F). 
The sodium hydroxide solution is evaporated to remove water 
and the molten mass is poured into iron barrels or into molds 
about the diameter of a lead pencil ; a flake form is also made. 
The chlorine is liquefied or used directly in manufacturing 
bleaching powder and other chlorine compounds. 

Sodium hydroxide is manufactured to some extent by adding 
lime to a boiling, dilute solution of sodium carbonate, but this 
process is being rapidly superseded by the electrolytic process. The 
essential chemical change is represented thus : — 



Ca(OH) 2 


+ 


Na 2 C0 3 = 


= 2 NaOH 


+ CaC0 3 


Calcium 




Sodium 


Sodium 


Calcium 


Hydroxide 




Carbonate 


Hydroxide 


Carbonate 



The solution of sodium hydroxide is treated as described above. 



426 INORGANIC CHEMISTRY 

Sodium Sulphate, Na 2 S0 4 , is one of the products obtained 
in the manufacture of sodium carbonate and of nitric acid 
(see above). At Stassfurt sodium sulphate is prepared by 
cooling to about 0° C. a mixture of magnesium sulphate 
and sodium chloride solutions; the equation is — 



MgS0 4 + 2 NaCl = 


Na 2 S0 4 + MgCl 2 


Magnesium Sodium 


Sodium Magnesium 


Sulphate Chloride 


Sulphate Chloride 



Sodium sulphate is a white solid. It dissolves readily 
in water, and when a concentrated solution made at 30° C. is 
cooled, large transparent crystals separate. They have the 
formula Na 2 SO 4 .10H 2 O and are called Glauber's salt 
from the discoverer. They lose water when exposed to air, 
and the salt continues to effloresce until it becomes an 
anhydrous powder. The crude salt is used in the glass and 
dyeing industries, and the purified salt as a medicine. 

Sodium Sulphite, Na 2 S0 3 , is a white solid prepared by 
passing sulphur dioxide into sodium hydroxide solution. 
The crystallized salt has the formula Na 2 S0 3 .7H 2 0. It 
interacts with sulphuric acid as follows : — 

+ H 2 



Na 2 S0 3 


+ 


H 2 S0 4 = 


=. Na 2 S0 4 + S0 2 


Sodium 






Sodium Sulphur 


Sulphite 






Sulphate Dioxide 



It is used as a source of sulphur dioxide and also as a pre- 
servative. 

Sodium Nitrate, NaNOs, is abundant in Chile, and is often 
called Chile saltpeter. It is a white solid, very soluble in 
water, especially hot water, and is slightly deliquescent. 
Large quantities are used as a fertilizer (either alone or mixed 
with compounds of potassium and of phosphorus) and for 
making nitric acid and potassium nitrate. 



SODIUM 427 

The natural deposits of sodium nitrate are in a dry region near 
the coast and cover over 200,000 acres. Chile controls the indus- 
try, and exports annually over a million tons. The crude nitrate, 
which is called caliche, is treated with water and then purified by 
crystallization into a product containing 94-98 per cent of sodium 
nitrate. The final mother liquid is a source of iodine. (See Iodine.) 

Sodium Nitrite, NaN0 2 , is a white solid. It is prepared 
by reducing sodium nitrate with lead, thus: — ■ 



NaN0 3 


+ 


Pb 


NaN0 2 


+ 


PbO 


Sodium 
Nitrate 




Lead 


Sodium 
Nitrite 




Lead 
Oxide 



It liberates the oxides of nitrogen (NO and N0 2 ) upon the 
addition of sulphuric acid, and is used extensively in 
manufacturing dyes. 

Sodium Dioxide or Peroxide, Na 2 3 , is a yellowish solid. 
It is used to bleach straw and delicate fabrics, and is an 
oxidizing agent. A fused form is sold as " oxone." With 
water it liberates oxygen, according to the equation — 

2 Na 2 2 + 2 H 2 = 2 + 4 NaOH 

Sodium Oxygen Sodium 

Dioxide Hydroxide 

Miscellaneous. — As stated above, sodium compounds 
impart an intense yellow color to a Bunsen flame. Most 
sodium compounds dissolve in water and they yield a 
colorless cation (Na + ). The atomic weight of sodium is 
23.00 and its valence is one. 

Sodium cyanide (NaCN) is used to extract gold from poor 
ores. The sodium phosphates, sodium thiosulphate, acid 
sodium sulphite, sodium silicate, and sodium tetraborate 
or borax have already been described. 



428 INORGANIC CHEMISTRY 

Potassium 

Occurrence. — This metal is not found free, but its com- 
pounds are abundant. The common minerals mica and 
feldspar are silicates containing potassium. By the decay 
of these and other minerals potassium compounds find their 
way into the soil, thence into plants and animals. Potas- 
sium salts are found in wood ashes, in suint (the oily sub- 
stance washed from sheep's wool), in beet-sugar residues, 
and in the deposits in wine casks. Sea water and mineral 
waters contain potassium salts, particularly potassium 
chloride and potassium sulphate. Many potassium salts 
are found at Stassfurt, and they are the source of most 
potassium compounds. About 2.5 per cent of the earth's 
crust is potassium. 

The Stassfurt deposits of the salts of potassium, magnesium, cal- 
cium, and sodium are near Magdeburg, Germany. The deposits are 
about five thousand feet thick and contain many salts which were 
deposited ages ago in beds or layers during the slow evaporation of an 
inclosed branch of the sea. The lowest bed is an enormous mass of 
more or less impure sodium chloride (called in this case rock salt or 
halite), which was deposited first. Upon this rest more or less regular 
layers of potassium and magnesium salts, higher still are calcium salts 
and at the top is a thick bed of sandstono. The different minerals 
are mined separately as far as possible and then separated into com- 
mercial products by an elaborate system of solution, evaporation, and 
crystallization. The most important Stassfurt potassium minerals 
are — 

Kainite . . . . . KC1, MgS0 4 . 3 H 2 0. 

Carnallite . . . . . . KC1, MgCl 2 . 6 H 2 0. 

Polyhalite .... K 2 S0 4 , MgS0 4 , 2 CaS0 4 . 2 H 2 0. 

Sylvite . . . . - . KC1. 

Picromerite .... K 2 S0 4 , MgS0 4 . 6 H 2 0. 

Many of the Stassfurt salts belong to a class of compounds 
called double salts, i.e. crystalline compounds of two or more 
normal salts with one another. Carnallite is a double 



POTASSIUM 429 

salt, and its formula is often written to emphasize this 
fact, thus KC1, MgCl 2 .6H 2 instead of KMgCl 3 .6H 2 0. 
Dilute aqueous solutions of double salts contain ions of the 
separate salts and exhibit no reactions which indicate com- 
bination of ions. There are other double salts. (See Alums.) 

Preparation. — Potassium is prepared by the electrolysis 
of fused potassium chloride. Formerly it was manufactured 
by heating to a high temperature a mixture of potassium 
carbonate and carbon, or of potassium hydroxide and iron 
carbide. 

Properties. — Like sodium, potassium is a soft, silver- 
white metal; it is light enough to float upon water — the 
specific gravity being .8621 (at 20° C). Its brilliant luster 
soon disappears in air, owing to rapid oxidation. Potassium 
as ordinarily seen is covered with a grayish coating, and like 
sodium must be kept under mineral oil. It melts at 62.5° 
C, and at a higher temperature burns with a violet-colored 
flame. This color is characteristic of burning potassium, 
and is a test for the metal and its compounds. Like sodium, 
it decomposes water at ordinary temperatures, though more 
energetically. The reaction corresponds to the equation — 



2K + 


2H 2 = 


= 2KOH 


+ H 2 


tassium 


Water 


Potassium 
Hydroxide 


Hydrogen 



The heat evolved immediately ignites the liberated hydrogen, 
and the melted potassium, surrounded by a violet flafrie, 
dashes to and fro upon the cold water. Potassium combines 
with the halogens and other elements more vigorously 
than sodium, and forms analogous compounds. 

Potassium Chloride, KC1, is found native as sylvite in 
the Stassfurt deposits. It is also obtained in large quan- 
tities by heating a concentrated solution of carnallite 



430 INORGANIC CHEMISTRY 

(KC1, MgCl 2 . 6 H 2 0) and allowing potassium chloride to 
crystallize out from the cool solution. It is a white solid 
which crystallizes in cubes and otherwise resembles sodium 
chloride. It is used chiefly to prepare other potassium salts, 
especially the nitrate, hydroxide, carbonate, and chlorate. 

Potassium bromide and iodide, which are analogous to the chloride, 
have been described (see these compounds). 

Potassium Nitrate, KN0 3 , is also called niter and salt- 
peter. It is formed in the soil near large cities in India 
and Persia by the decomposition of nitrogenous organic 
matter in the presence of potassium salts. (See Nitrification.) 
It is manufactured by mixing hot, concentrated solutions 
of native sodium nitrate and potassium chloride, which 
interact thus : — 



NaN0 3 


+ KCl = 


kno 3 


+ 


NaCl 


Sodium 
Nitrate 


Potassium 
Chloride 


Potassium 
Nitrate 




Sodium 
Chloride 



The sodium chloride, being the less soluble of the two, pre- 
cipitates as the solution cools, and is removed. By further 
evaporation the potassium nitrate (together with a little 
sodium chloride) crystallizes out as " niter meal." 

Potassium nitrate is a white solid. It dissolves only to 
a slight extent in cold water, but very freely in hot water. 
It is not deliquescent. If a solution is cooled slowly, large 
prismatic crystals are formed w r hich do not contain water 
of crystallization. Such crystals have cavities in which 
there is impure water; this water cannot be removed by 
drying the crystals. When heated, these crystals decrepi- 
tate like those of sodium chloride. The taste is salty and 
cooling. It melts at about 339° C, and on further heating 
changes into potassium nitrite (KN0 2 ) and oxygen. It 
is stable in the air at ordinary temperatures, but at a high 
temperature potassium nitrate gives up oxygen readily, 



POTASSIUM 431 

especially to charcoal, sulphur, and organic matter. This 
oxidizing power leads to its extensive use in making gun* 
powder, fireworks, and matches. 

Gunpowder is a mixture of potassium nitrate, charcoal, and sulphur. 
The ingredients are first purified, pulverized, and thoroughly mixed. 
This mixture is pressed, while damp, into a thin sheet; and the press 
cake thus formed is broken into small grains, which are sorted by 
sieves. The grains are then smoothed or glazed by rolling them in a 
barrel, again sifted, and finally dried at a low temperature. The pro- 
portions differ with the use of the powder. Modern black powder 
contains about 75 per cent of potassium nitrate, 15 of charcoal, and 
10 of sulphur. When gunpowder burns in a closed space, a large 
volume of gas is suddenly formed. So enormously is this gas expanded 
by the heat that in an open space it would fill several hundred times 
the space taken by the powder itself. The pressure exerted by this 
expanding gas is many tons to the square inch. It is this pressure 
which forces the charge from a gun and tears a rock to pieces. The 
chemical changes attending the explosion of confined gunpowder are 
complex, as may be seen by the following equation: — 

I6KNO3 + 2IC + 5S = 5 K2CO3 + 13 CO2 + K 2 S0 4 + 3 CO 

Potassium Carbon Sulphur Potassium Carbon Potassium Carbon 
Nitrate Carbonate Dioxide Sulphate Monoxide 

+ 2K2S2 + 8N2 

Potassium Nitrogen 
Sulphide 

The equation for the explosion when unconfined is much simpler, 
thus : — 

2 KNO3 + 3 C + S = 3 C0 2 + N 2 + K 2 S 

Gunpowder is being rapidly replaced by smokeless powders. (See 
Cellulose.) 

Potassium Chlorate, KC10 3 , is a white, crystalline, lustrous 
solid. It tastes like potassium nitrate. It melts at about 
370° C, and at a high temperature decomposes into oxygen 
and potassium chloride as final products, thus: — - 

KCIO3 = KC1 + 3 

Potassium Potassium Oxygen 
Chlorate Chloride 



432 INORGANIC CHEMISTRY 

It is used to prepare oxygen, and in the manufacture of dyes, 
matches, and fireworks. In the form of " chlorate of potash 
tablets " it is used as a remedy for sore throat. 

Potassium chlorate was formerly manufactured by passing 
chlorine into calcium hydroxide (milk of lime) and adding 
potassium chloride to the mixture. The simplest equations 
for the complex changes may be written thus: — 

6 Ca(OH) 2 + 6 Cl 2 = Ca(C10 3 ) 2 + 5 CaCl 2 + 6 H 2 

Calcium Chlorine Calcium Calcium 

Hydroxide Chlorate Chloride 

Ca(C10 3 ) 2 + 2KC1 = 2KC10 S + CaCl 2 

Potassium Potassium 
Chloride Chlorate 

This salt is now made by the electrolysis of a hot, con- 
centrated solution of potassium chloride. The two products 
— chlorine and potassium hydroxide — interact thus : — 

3 Cl 2 + 6 KOH = KC10 3 + 5 KC1 + 3 H 2 

Chlorine Potassium Potassium Potassium Water 

Hydroxide Chlorate Chloride 

Potassium Perchlorate (KC10 4 ) is a white crystalline solid formed by 
heating potassium chlorate. 

Potassium Carbonate, K2CO3, is a white solid. It is 
very soluble in water and deliquesces when exposed to the 
air, becoming a thick liquid at first and finally a solid. The 
property of deliquescence is due to the fact that the vapor 
pressure of its saturated solution at ordinary temperatures 
is less than the average pressure of the water vapor in the 
air. The solid residue is potassium bicarbonate, w T hich is 
formed by the slow absorption of carbon dioxide from the 
air. A solution of potassium carbonate has a marked 
alkaline reaction. (See Sodium Carbonate.) It was formerly 
prepared in large quantities by extracting wood ashes with 
water and evaporating the solution to dryness; this process 
is still employed in some localities. The crude salt thus 
obtained has long been called potash and a purer product 



POTASSIUM 433 

is known as pearlash. The term potash is also sometimes 
applied to potassium hydroxide (KOH) and to potassium 
oxide (K 2 0). It is used in the manufacture of hard glass, 
soft soap, and potassium compounds. 

Potassium carbonate is obtained to some extent from suint by 
igniting the greasy mass and extracting the potassium carbonate with 
water. Beet-sugar residues also furnish potassium carbonate. After 
the sugar has been obtained from the beet sirup, the molasses is 
changed by fermentation into alcohol, which is distilled off; the 
liquid residue is evaporated to dryness and ignited, and the potassium 
carbonate extracted with water. Potassium carbonate is prepared 
by igniting the crude cream of tartar collected from the deposits in 
wine casks, and for this reason it is sometimes called salt of tartar. 
All these sources emphasize the intimate relation of potassium 
compounds to vegetable and animal life. Much of the commercial 
potassium carbonate is now made from potassium chloride by the 
Leblanc process. Another process is used, principally in Germany, 
owing to the abundance of crude potassium salts at Stassfurt. It 
consists essentially in forcing carbon dioxide into a solution of po- 
tassium chloride containing suspended magnesium carbonate, de- 
composing the complex compound of potassium and magnesium 
(MgC03 . HKCO3 . 4 H 2 0) with steam, and evaporating the filtered 
solution of potassium carbonate. 

The name potassium comes from the word potash. The symbol, 
K, is from kalium, the Latin equivalent of kali, which is derived from 
an Arabic term for an alkaline substance. 

Potassium Hydroxide or Caustic Potash, KOH, is a white 
brittle solid, resembling sodium hydroxide. It absorbs 
water and carbon dioxide very readily; and if exposed to the 
air, it soon becomes a thick solution of potassium carbonate. 
Like sodium hydroxide, it dissolves in water with evolution 
of heat, forming a markedly alkaline, caustic solution. 
Besides its use in the laboratory, large quantities are con- 
sumed in making 50ft soap. 

Potassium hydroxide was formerly made in the same way 
as sodium hydroxide, viz. by adding lime or milk of lime to 



434 INORGANIC CHEMISTRY 

a boiling dilute solution of potassium carbonate, the equation 
for the change being — 



Ca(OH) 2 


+ K 2 C0 3 = 


= 2KOH 


+ 


CaC0 3 


Milk 


Potassium 


Potassium 




Calcium 


of Lime 


Carbonate 


Hydroxide 




Carbonate 



It is now manufactured by the electrolysis of a solution of 
potassium chloride, the process being like the one used for 
sodium hydroxide. 

Other Potassium Compounds. — Potassium Cyanide (KCN) 
is a white solid, very poisonous, very soluble in water, and 
has an odor like bitter almonds. (See Cyanogen, Chapter 
XVII.) It is used in preparing electroplating solutions and 
in extracting gold from poor ores. Potassium Sulphate 
(K2S0 4 ) is manufactured from kainite, and is largely used 
as a fertilizer and in making potassium compounds. 

Relation of Potassium to Life. — Potassium, like nitrogen 
and phosphorus, is essential to the life of plants and animals. 
The ash of many common grains, vegetables, and fruits con- 
tains potassium as the carbonate. Potassium salts are 
supposed to assist in the formation of starch, just as phos- 
phorus is indispensable to the transformation of nitrogen 
compounds. Potassium salts taken from the soil by plants 
must be returned, if the soil is to be productive. Sometimes 
crude kainite is used extensively as a fertilizer; but wood 
ashes, or the sulphate and chloride, are often used to supply 
potassium salts. (See Relation of Phosphorus to Life, 
Chapter XXIII.) 

Miscellaneous. — As already stated, potassium compounds 
impart a delicate violet tint to a Bunsen flame. Most potas- 
sium compounds are soluble in water, and such solutions 
contain colorless potassium ions (K + ). The atomic weight 
of potassium is 39,10 and the valence is one. 



AMMONIUM COMPOUNDS 435 

Lithium, Li, is a silver-white metal and has the specific gravity 
of only .534 (at 20° C), being the lightest of the metallic elements. 
Its compounds are widely distributed in small quantities in minerals, 
mineral waters, and plants. Lithia water and citrate of lithium are 
often prescribed as a remedy for diseases of the kidneys. Lithium 
compounds color the Bunsen flame a bright red — a delicate test 
for the element. 

Rubidium, Rb, and Caesium, Cs, have properties and form corn- 
pounds analogous to those of potassium. 

The Alkali Metals, as already stated, form a natural family. The 
properties of the metals are quite similar, and the chemical activity 
increases in passing from lithium (at. wt. 6.94) to caesium (at. wt. 
132.81). All decompose water, yielding hydrogen and an hydroxide 
of the metal. The hydroxides are active bases, and the familiar ones 
long ago gave the name alkali to the family. Analogous compounds 
are very much alike ; indeed, in many operations, it makes little dif- 
ference whether sodium or potassium compounds are used, though the 
former are usually preferred on account of their lower price. 



Ammonium Compounds 

Introduction. — We found in Chapter XIII that ammonium 
(NH 4 ) is a metallic radical, i.e. a group of elements which 
acts like an atom of a metal in chemical changes. Its most 
familiar compound is ammonium hydroxide (NH 4 OH), 
which has the properties of a base and resembles sodium 
and potassium hydroxides. Other compounds of ammonium, 
especially certain salts, are analogous to the corresponding 
salts of sodium and potassium. Hence, ammonium com- 
pounds, except ammonium hydroxide, are appropriately 
described in this chapter. 

Ammonium Chloride, NH 4 C1, is prepared by passing 
ammonia gas into dilute hydrochloric acid, by mixing am- 
monium hydroxide and hydrochloric acid, or by letting the 



436 INORGANIC CHEMISTRY 

two gases mingle. The equation for the essential reaction 
is — 

NH S + HC1 = NH 4 C1 

Ammonia Hydrochloric Ammonium 

Acid Chloride 

It is convenient to regard this compound as the ammonium 
salt of hydrochloric acid, as if it were formed by replacing 
the hydrogen of the acid by ammonium, just as sodium forms 
sodium chloride. 

Ammonium chloride is a white, granular, fibrous, or crys- 
talline solid, with a sharp, salty taste. It dissolves easily 
in water, and in so doing lowers the temperature markedly. 
When ammonium chloride is heated to a high temperature 
(about 350° C), it volatilizes and dissociates into ammonia 
and hydrogen chloride; these gases reunite to form ammo- 
nium chloride as the temperature falls. 

Large quantities of ammonium chloride are made at one 
stage of the manufacture of ammonium hydroxide by passing 
the gas into hydrochloric acid. The crude product is called 
" muriate of ammonia" to indicate its relation to muriatic 
(or hydrochloric) acid. It is largely used in charging 
Leclanche batteries, as an ingredient of soldering fluids, in 
galvanizing iron, and in textile industries. The crude salt 
is purified by heating it gently in a large iron or earthen- 
ware pot, with a dome-shaped cover; the ammonium 
chloride volatilizes easily and then crystallizes in the pure 
state as a fibrous mass on the inside of the cover, but the 
impurities remain behind in the vessel. The process of 
vaporizing a solid substance and then condensing the vapor 
directly into the solid state is called sublimation. It differs 
from distillation in that the substance does not pass through 
an intermediate liquid state. The product of sublimation 
is called a sublimate. Sublimed ammonium chloride is 
known as sal ammoniac. 



AMMONIUM COMPOUNDS 437 

Ammonium Sulphate, (NH 4 ) 2 S0 4 , is made by passing 
ammonia gas into sulphuric acid, or by adding ammonium 
hydroxide to the acid, thus: — 

2NH 4 OH + H 2 S0 4 = (NH 4 ) 2 S0 4 + 2H 2 

Ammonium Ammonium 

Hydroxide Sulphate 

The commercial salt is a grayish or yellowish solid. It is 
used as a component of fertilizers, since it is rich in nitrogen, 
and in making ammonium alum and other ammonium com- 
pounds. 

Ammonium Sulphide, (NH 4 ) 2 S, is prepared by passing 
hydrogen sulphide gas into ammonium hydroxide, thus : — 

H 2 S + 2NH 4 OH = (NH 4 ) 2 S + 2H 2 

Hydrogen Ammonium Ammonium 

Sulphide Hydroxide Sulphide 

The normal sulphide is unstable and forms acid ammonium 
sulphide (also called ammonium hydrosulphide, NH 4 HS). 
Hence, a solution of ammonium sulphide, as usually prepared, 
contains both the normal and acid sulphide; it smells of 
hydrogen sulphide and ammonia, and if exposed to the air, 
it ultimately changes into ammonia, sulphur, and water. 
A solution of ammonium sulphide is used in qualitative 
analysis to precipitate as sulphides certain metals of the 
third group, e.g. nickel, cobalt, zinc, and manganese. Am- 
monium sulphide solutions dissolve sulphur, thereby forming 
a solution of complex sulphides called yellow ammonium 
sulphide or ammonium polysulphide. The polysulphide is 
used in qualitative analysis. (See Test for Arsenic.) 

Ammonium Nitrate, NH 4 N0 3 , is made by passing ammonia 
into nitric acid, or by allowing ammonia gas and the vapoi 
of nitric acid to mingle, thus : — 

NH 3 + HN0 3 = NH 4 NO s 

Ammonia Nitric Ammonium 

Acid Nitrate 



438 INORGANIC CHEMISTRY 

It is a white salt which forms beautiful crystals. It dis- 
solves easily in water with a fall of temperature. When 
gently heated it decomposes into nitrous oxide (N 2 0) and 
water, and its chief use is in the preparation of nitrous 
oxide (see this compound). 

Ammonium Carbonate, (NH 4 ) 2 C0 3 , as usually found in 
commerce is a mixture of acid ammonium carbonate 
(HNH4CO3) and a related compound. It is a white solid ; 
on exposure to air it loses ammonia and forms the acid car- 
bonate. It is used to prepare some kinds of baking pow- 
der, to scour wool, as a medicine, and to prepare smelling 
salts, since it gives off ammonia readily. The solution is 
used in qualitative analysis to precipitate barium, stron- 
tium, and calcium. 

Miscellaneous. — Most ammonium compounds dissolve 
in water, and such solutions contain colorless ammonium 
ions (NH 4 + ). Attempts to isolate the radical ammonium 
have thus far failed. The valence of ammonium is one. 

Ammonium hydroxide and sodium ammonium phosphate 
(microcosmic salt) have already been considered. 

Spectrum Analysis 

When sodium or one of its compounds is introduced into 
a Bunsen flame, a vivid yellow color is imparted to the flame; 
the color yielded by potassium or its compounds is a delicate 
violet. Many elements, especially the metals, behave 
similarily. These colors, as already stated, often serve as 
a simple test for the element (free or combined). But if 
only a minute amount of the substance is available or an 
intense color masks a faint one, the test fails or is unreliable; 
in a few cases, too, the colors are much alike. Nevertheless, 
it is possible to detect elements in a flame even though the 



SPECTRUM ANALYSIS 



439 



color is faint or obscured. This is accomplished by the 
spectroscope (Fig. 67). This instrument consists essentially 




Fig. 67. — Diagram of a spectroscope showing collimator, slit (£), tele- 
scope, and lenses (L 1; L 2 , L s ). 

of three parts, (a) a tube called the collimator containing 
a narrow slit at one end through which the light enters, (6) 
a triangular glass prism (placed with its edges parallel to 
the slit) through which the light passes as it comes from 
the collimator, and (c) a telescope located at such an angle 
that the light can be viewed as it emerges from the prism. 
When ordinary .white light, which consists of rays of all 
colors, enters the slit and falls upon the prism, the rays of 
light in passing through the 
prism are bent. That is, 
the numberless rays making 
up the white light emerge 
at an angle to the line 
along which they entered, FlG 68 
the red being bent the least, 
the violet the most (Fig. 68). Consequently, the emergent 
light, if caught upon a piece of ground glass or viewed 
through the telescope, is no longer white, but is a continuous 
band of colors arranged like the familiar colors of the rain- 
bow. This band of colors consists of a series of colored 
images of the slit and is called a spectrum. The production 
and examination of a spectrum is termed spectrum analysis. 




- Dispersion of light by a prism. 



440 INORGANIC CHEMISTRY 

White light contains all the colors of the spectrum. But 
when colored light, such as that from a sodium or potassium 
flame, is examined by the spectroscope, instead of the con- 
tinuous band of colors, we see only those images of the slit 
which correspond to the rays of light in the sodium or 
potassium flame. Thus, the sodium flame gives one image 
of the slit (under usual conditions), which is seen as a bril- 
liant yellow vertical line, besides other and minor ones 
against a black background; similarly, potassium gives 
two conspicuous lines, a red and a violet. Each element, if 
heated to a sufficiently high temperature, has its own series 
of colored lines, which is called its line spectrum. The 
spectrum of some elements is complex, though many have 
certain lines which are so conspicuous that the element can 
be readily detected in a mixture. 

In the laboratory the spectroscope is used to detect the 
presence of certain elements, more especially the metals of 
the alkali and alkaline earth families. If a small piece of 
the metal or one of its compounds (preferably the chloride) 
is put on a platinum wire and held in the Bunsen flame 
before the slit, the characteristic line spectrum of the ele- 
ment can be readily recognized by looking into the telescope. 
Several elements can be detected in a mixture, for although 
certain lines may coincide or overlap, other lines are con- 
spicuous enough to reveal the presence of the components. 
Minute quantities are detected by the spectroscope, since 
the light, though too faint to affect the eye, is concentrated 
by the spectroscope into a bright line (or lines), which stands 
out against the black background. Consequently rare 
elements, which can be obtained only in small quantities ox* 
with great difficulty, are usually studied spectroscopically. 
Thus, Bunsen, who (with Kirchhoff) devised the improved 
spectroscope, discovered and studied the rare metals rubid- 
ium and caesium. Mention has already been made of the 



PROBLEMS AND EXERCISES 441 

fact that during the last few years the spectroscope has 
been especially serviceable in studying argon and helium 
and the related gases. 

Problems and Exercises 

1. How many pounds of sodium could be made (theoretically) 
from 20 metric tons of sodium hydroxide? 

2. What is the weight of a cubic meter of potassium ? (Assume 
a cubic centimeter of H 2 to weigh 1 gm. at 20° C.) 

3. How much sulphuric acid is needed to convert 10 gm. of 
sodium chloride into sodium sulphate? 

4. What volume of carbon dioxide (at standard conditions) will 
be formed by heating 72 gm. of sodium bicarbonate? 

5. What weights of potassium and water are needed to produce 
50 1. of hydrogen (standard conditions) ? 

6. What weight of sulphur is needed to convert 80 gm. of sodium 
sulphite into sodium thiosulphate ? (Equation is Na 2 S0 3 + S = 
Na 2 S 2 3 .) 

7. What weight of Na^COs . 10 H 2 will 2000 lb. of NaCl pro- 
duce? 

8. Suppose 10 gm. of gunpowder, when exploded, yielded 3 1. of 
gas measured at 0° C. and 760 mm. What would be the volume at 
1800°? What pressure would be exerted if the volume was kept 
unchanged ? 

9. Calculate the simplest formulas corresponding to (a) Na = 
32.39, S = 22.54, = 45.07; (6)Na = 36.5, S = 25.4, = 38.09; 
(c) Na = 29.11, S = 40.50, = 30.38. 

10. Starting with sodium, how would you prepare successively 
the chloride, sulphate, sulphide, carbonate, chloride, hydroxide, 
and metal ? Also from KC1 the following in succession : KN0 3 , 
HN0 3 , NaN0 3 , HNO3, KNO3, KNO2? Also from sodium, its 
oxide, hydroxide, chloride, acid sulphate, normal sulphate? 

11. Write the formulas of the following compounds by applying 
the principle of valence (see Chapter XIV) : Sodium chlorate, so- 
dium acetate, sodium fluoride, sodium phosphate (ortho), potassium 
manganate, acid potassium sulphite, lithium carbonate, lithium 
chloride, lithium phosphate (ortho). 

12. Indicate the ions which are in dilute aqueous solutions of 
the following : Potassium nitrate, ammonium chloride, picromerite. 



CHAPTER XXVI 
Copper — Silver — Gold 

Introduction. — These metals are related and constitute a 
family in Group I of the periodic classification/ but they do 
not have such marked family characteristics as the alkali 
metals. The metals, as well as their alloys and compounds, 
have many domestic and commercial uses. 

Copper 

Copper has been known for ages. Domestic utensils and 
weapons of war containing copper were used before similar 
objects of iron. The Romans obtained copper from the 
island of Cyprus. They called it cuprium aes (i.e. Cyprian 
brass), which finally became simply cuprum. From cuprum 
we obtain the symbol Cu and the terms cuprous and cupric. 

Occurrence of Copper. — Copper, both free and combined, 
is an abundant element. Single masses of native or metallic 
copper weighing many tons are found in the Michigan 
mines in the Lake Superior region. Besides the native 
copper, the most valuable ores are copper sulphide (chal- 
cocite or copper glance, Cu 2 S), copper oxide (cuprite or 
ruby ore, Cu 2 0), the copper-iron sulphides (copper pyrites 
or chalcopyrite, CuFeS 2 , and bornite, Cu 3 FeS 3 ), and the 
basic carbonates (malachite, CuC0 3 . Cu(OH) 2 , and azurite, 
2CuC0 3 .Cu(OH) 2 ). 

Native copper comes chiefly from Michigan, the copper- 
iron sulphide ores from Montana, and the oxide and car- 
bonates from Arizona. 

442 



COPPER 443 

Metallurgy of Copper. — Copper is extracted from its ores 
by processes which vary with the composition of the ore. 
(1) Native copper ore is first crushed, then washed to re- 
move rocky impurities, and the concentrated product 
finally melted. (2) The carbonates and oxide are reduced 
by roasting them with coke in a blast furnace. The general 
chemical change is reduction and may be represented thus: — 



Cu 2 


+ c = 


: 2CU 


+ CO 


Copper 
Oxide 


Carbon 


Copper 


Carbon 
Monoxide 



(3) The smelting of copper-iron sulphides is a complicated 
process. The ore is crushed and washed, and then roasted 
in a furnace. This operation removes the adhering rock 
and changes much of the sulphide into an oxide. The 
roasted mass is next heated with coal and sand in a shaft 
or a reverberatory furnace, whereby the iron is largely 
changed into a fusible silicate which runs off as a slag. 
The remaining " matte," as it is called, contains from 35 to 
50 per cent of copper besides iron, sulphur, and arsenic 
(as well as gold and silver), and is further treated, (a) It 
is roasted and melted until all the iron and arsenic are re- 
moved and mainly copper sulphide remains. This is finally 
roasted to convert it partly into an oxide, and the mixture 
of sulphide and oxide is again melted ; the sulphur passes 
off as sulphur dioxide, and the copper is left behind. The 
equation for this final change is — 



2CuO 


+ 


Cu 2 S = 


= 4Cu 


+ so 2 


Copper 




Copper 


Copper 


Sulphur 


Oxide 




Sulphide 




Dioxide 



(b) As a rule the process is hastened by pouring the molten 
" matte " into a silica-lined converter (see Fig. 81) and 
blowing air through the molten mass. The sulphur passes 
off as sulphur dioxide and the iron forms a fusible slag. The 



444 



INORGANIC CHEMISTRY 



t 


r— 


of 


1 

[ 


C r 


4 


^r 1 

A 


1 


A«^m^^^ 



product is called blister copper and is about 98 per cent 
pure. It is cast into thick plates called anodes and purified 
by electrolysis. 

Purification of Copper. — The anodes are connected with 
the positive electrode of a powerful battery or dynamo and 
suspended in a solution of copper sulphate and sulphuric 
acid. Sheets of pure copper are made the cathodes and 
dip into the solution as shown in Figure 69. When the cur- 

rent passes, pure copper 

_ j- leaves the anodes and be- 
comes deposited upon the 
cathodes; the impurities 
either remain in solution 
or fall to the bottom of 
the tank as a slime, from 
which gold and silver are 
extracted in appreciable 
quantities. The operation 
can be interpreted by prin- 
ciples already discussed. The copper ions (Cu ++ ) migrate to 
the cathode, where they lose their electric charges and are 
deposited as metallic copper (Cu). The S0 4 -ions (S0 4 ~~) 
migrate to the anode, where they likewise lose their charges, 
become ordinary chemical S0 4 -groups, and unite with copper 
from the anode to form CuS0 4 , which, however, immediately 
ionizes into Cu ++ and S0 4 ~". The gold and silver do not 
ionize and the zinc (if present) remains in the solvent as 
ionic zinc (Zn ++ ) . Thus the solution is constantly supplied 
with ions of copper and S0 4 , and pure copper is removed 
from the anodes and deposited in equivalent amounts 
upon the cathodes. New anodes and cathodes are supplied 
as needed. Electrolytic copper, a§ it is often called, is 
exceedingly pure. 



Fig. 69. — Apparatus for the prepara- 
tion of pure copper by electrolysis. 
A, A, A are anodes, and C, C, C are 
cathodes. 



COPPER 445 

Properties of Copper. — Copper is a bright metal, distin- 
guishable from all others by its peculiar reddish color. It 
is flexible, hard, and tough; its malleability and ductility 
allow it to be drawn out into wire and rolled into very thin 
sheets. Its specific gravity varies slightly with the method 
of treatment, but is about 8.9. The melting point is about 
1083° C. Copper is an excellent conductor of both heat 
and electricity. In dry air, it turns dull; and in moist air 
it gradually becomes coated with a greenish basic copper 
carbonate. Heated in the air, it is changed into the black 
copper oxide, and at a high temperature it colors a flame 
emerald green. Copper does not liberate hydrogen readily 
from dilute acids. With nitric acid it forms copper nitrate 
and oxides of nitrogen (see Oxides of Nitrogen); with hot 
concentrated sulphuric acid it yields copper sulphate and 
sulphur dioxide (see Sulphur Dioxide). Hydrochloric acid 
has little effect upon it. Copper replaces certain metals if 
suspended in solutions of their compounds, e.g. a clean 
copper wire soon becomes coated with mercury, if placed 
in a solution of any mercury compound; on the other hand, 
certain metals, such as iron, zinc, and magnesium, remove 
copper from its solution, e.g. a nail or knife blade soon be- 
comes coated with copper if clipped into a solution of any 
copper compound. Scrap iron is sometimes used to pre- 
cipitate copper on a large scale. (See Displacement of 
Metals below.) 

Uses of Copper. — Next to iron, copper is the most useful 
metal. Enormous quantities of wire are used in operating 
the telegraph, cable, telephone, electric railway, and electric 
light. Sheet copper is made into household utensils, boilers, 
and stills. Copper bolts, nails, rivets and sheathing are 
used on ships, because copper is only slowly corroded by 
moist air and salt water. All nations use copper as the 



446 INORGANIC CHEMISTRY 

chief ingredient of small coins. Electrical apparatus utilizes 
much copper. Maps, etchings, and some kinds of engrav- 
ings are printed from copper plates; calico is printed from 
a copper cylinder upon which the design is engraved. Books 
are printed and illustrated from electrotypes, made by de- 
positing a film of copper upon an impression of the type 
or design in wax. In a similar way many objects are cop- 
per plated. (See Electrotyping and Electroplating, Chap- 
ter XI.) 

Alloys of Copper are important. Brass is a bright yellow 
alloy containing 63 to 72 per cent of copper, the remainder 
being zinc. It is made by melting these metals together. 
It can be drawn into wire, hammered into any shape, and 
turned in a lathe. It is harder than copper, and melts at a 
lower temperature. On account of its durability, elasticity, 
and other properties, it has many uses for which copper 
and zinc are not suited. Pinchbeck, Muntz metal, Bath 
metal, Dutch metal (leaf or "gold")* are varieties of brass. 
Muntz metal is often used in place of sheet copper as sheath- 
ing for the bottoms of ships, because it becomes corroded 
very slowly. Typical bronze contains different proportions 
of copper, zinc, and tin; some antique bronzes contain lead 
or iron. The per cent of copper is 70 to 95, of zinc 1 to 25, 
of tin 1 to 18. The proportions in the British bronze coinage 
are copper 95, zinc 1, tin 4. On account of its fusibility, 
beautiful color, and extreme durability, bronze is used for 
statues, memorial tablets, coins, and medals. The ancients 
made it into weapons of war and household utensils. Cannon 
were formerly made of bronze, but for this purpose steel is 
now used. Phosphor bronze contains a small per cent of 
phosphorus and tin, and manganese bronze about 30 per 
cent of manganese; they are tougher than ordinary bronze, 
and are used to make steamship propellers and parts of 



COPPER 447 

machines. Silicon bronze is copper with traces of iron and 
silicon; its tenacity makes it especially serviceable for tele- 
graph, trolley, and telephone wires. Aluminium bronze 
contains 90 to 95 per cent copper, the rest being aluminium. 
It is a hard, yellow-brown, elastic alloy, and is used in con- 
structing hulls of yachts; its lightness, strength, and re- 
sistance to chemicals adapt it to many other uses. 

Gun metal is about 90 per cent copper and 10 per cent zinc; it was 
formerly used in making cannon, and is now used to some extent in 
making firearms. Bell metal contains about 75 per cent copper and 
25 per cent zinc. Speculum metal contains about 70 per cent copper, 
30 per cent tin, and traces of zinc, nickel, and iron; it takes a brilliant 
polish and is used in optical instruments. The numerous varieties 
of German silver contain different proportions of copper, nickel, and 
zinc. The per cent of copper is 50 to 60, of nickel 20 to 25, and of 
zinc about 20. In color and luster it resembles silver, for which it is 
often substituted. Its power to conduct electricity is only slightly 
affected by changes of temperature, hence it is often used in resistance 
coils. Chinese Pakfong (or paktong) is a variety of German silver. 
The nickel coins of Germany and the United States contain 75 per 
cent copper and 25 per cent nickel. Copper is also a constituent of 
many other coins. Britannia metal and white metal, in which copper 
is the minor constituent, are described under Alloys of Tin. 

Compounds of Copper. — Copper forms two series of com- 
pounds, the cuprous and the cupric. Thus, there are 
cuprous oxide (CU2O) and cupric oxide (CuO), cuprous 
chloride (CuCl) and cupric chloride (CuCl 2 ). The cuprous 
compounds contain a larger proportion of copper than the 
corresponding cupric compounds. The valence of copper is 
one in the cuprous series and two in the cupric. Not every 
member of each series, however, is important or even well 
known. Solutions of cupric salts contain blue cupric ions 
(Cu ++ ) and those of cuprous salts colorless cuprous ions 
(Cu + ). Copper compounds are poisonous. Cooking uten- 
sils made of copper should be used with care; vege- 



448 INORGANIC CHEMISTRY 

tables, acid fruits, and fruit preserves, if boiled in such 
vessels, should be removed as soon as cooked. The vessels 
themselves should be kept bright to prevent the formation 
of soluble copper salts which might contaminate the contents. 

Cuprous Oxide, Cu 2 0, occurs native as cuprite or ruby 
ore. It can be obtained as reddish powder by heating a 
mixture of solutions of copper sulphate, Rochelle salt 
(potassium sodium tartrate), sodium hydroxide, and glucose. 
This oxide colors glass ruby red. It is a beautiful mineral 
and a valuable ore. 

Cupric Oxide, CuO, is a black solid formed by heating 
copper in a current of oxygen or by heating copper nitrate 
or other cupric salts intensely. It is reduced to metallic 
copper when heated in a current of hydrogen or with sub- 
stances containing hydrogen or carbon, thus : — 



CuO 


+ H 2 = 


Cu 


+ 


H 2 


Copper 
Oxide 


Hydrogen 


Copper 




Water 



This property has led to its use in determining the amount 
of hydrogen or carbon in compounds. (See also Gravimetric 
Composition of Water.) 

Copper Sulphate or Cupric Sulphate, CuS0 4 , is the most 
useful compound of copper. It is a blue, crystalline solid, 
and is often called blue vitriol or bluestone. The crystal- 
lized salt (CuS0 4 . 5 H 2 0) loses water slowly in the air; heated 
to 240° C, all the water escapes, leaving a whitish powder. 
This anhydrous copper sulphate, as it is often called, absorbs 
water from alcohol and similar liquids ; and when added to 
water, it again becomes blue. An aqueous solution of cop- 
per sulphate has an acid reaction. As already stated, this 
is due to hydrolysis. The cupric ions (Cu ++ ) combine with 



COPPER 449 

hydroxyl ions (from the slightly dissociated water) to form 
the slightly dissociated cupric hydroxide (Cu(OH) 2 ). This 
removal of OH-ions, although slight, leaves in the solution 
enough hydrogen ions (H + ) to turn blue litmus to red. (See 
Hydrolysis, Chapter X.) 

Copper sulphate is used in electric batteries, in making 
other copper salts, in calico printing, dyeing, copper plating, 
in preserving timber; recently very dilute solutions have 
been used to destroy certain forms of objectionable organic 
matter in drinking water. It is a germicide and is one in- 
gredient of certain mixtures (such as Bordeaux mixture) 
which are sprayed upon trees to destroy fungi and kill in- 
sects. 

Copper sulphate is prepared on a large scale by treating 
copper with dilute sulphuric acid, the equation for the 
chemical change being — 

2Cu + 2H 2 S0 4 + 2 = 2CuS0 4 + 2H 2 

Copper Sulphuric Oxygen Copper Water 

Acid Sulphate 

A large proportion of the copper sulphate of commerce is 
obtained as a by-product in refining gold and silver with 
sulphuric acid (see below). 

Copper Nitrate, Cu(N0 3 ) 2 , is a blue, crystalline solid, 
formed by the interaction of dilute nitric acid and copper 
or copper oxide. It is a cupric salt. It is deliquescent, 
very soluble in water, and is readily decomposed by heat 
into cupric oxide and oxides of nitrogen. 

Cuprous Sulphide, Cu 2 S, is the bluish black mineral chalcocite. 
Cupric sulphide, CuS, is the black precipitate formed by the inter- 
action of hydrogen sulphide and a cupric salt. 

Copper Carbonates. — Malachite (CuCOs . Cu(OH) 2 ) and azurite 
(2 CUCO3 . Cu (OH) 2 ) are basic carbonates. Both, occur as minerals. 



450 INORGANIC CHEMISTRY 

malachite being bright green and azurite a magnificent blue. They are 
valuable ores of copper. Malachite is easily polished and is used as 
an ornamental stone for pillars, mosaics, and table tops. 

Copper Acetates. — Verdigris is copper acetate (CU3 (OH)2 . (C2H302)4), 
which is used to some extent in making gi;een paint and Paris green. 
The latter is copper aceto-arsenite and is used to exterminate potato 
bugs and other insects. Verdigris is a basic salt. 

Copper-Ammonia Compounds. — When a little ammonium hydrox- 
ide is added to the solution of a copper salt, a whitish, gelatinous 
precipitate is formed which upon the addition of an excess of am- 
monium hydroxide becomes a deep blue solution. If cupric sulphate 
is used, complex compounds can be obtained from this solution, e.g. 
Cu(NH 3 )4S0 4 . The blue color of the solution is due to the com- 
plex ion Cu(NH 3 ) 



++ 



Tests for Copper. — (1) The reddish color, peculiar "cop- 
pery" taste, and green color imparted to a Bunsen flame 
serve to identify metallic copper. (2) An excess of ammonium 
hydroxide added to the solution of a copper compound 
produces a deep blue solution. (3) A few drops of acetic 
acid and of potassium ferrocyanide solution added to a 
dilute solution of a copper compound produce a brown 
gelatinous precipitate of cupric ferrocyanide (Cu 2 Fe(CN) 6 ). 

Displacement of Metals. — The deposition of metallic 
copper when certain metals are put into solutions of cop- 
per salts and the displacement of mercury from solutions 
of its salts by metallic copper itself are examples of a kind 
of chemical change in which most metals can participate. 
Experiment shows that familiar metals can be arranged in a 
series (see page 451) based on their displacing power. In 
this series each free metal displaces the succeeding metals 
from their solutions, and is in turn displaced from solution 
by those metals which precede. 



COPPER 



451 



Electrochemical or Electromotive Series of Metals 



Magnesium + 1.31 


Cobalt 


+ 0.05 


Copper - 0.58 


Aluminium -f- 1.04 


Nickel 


-0.02 


Mercury — 0.99 


Zinc + 0.52 


Lead 


-0.12 


Silver - 1.04 


Iron +0.19 


Tin 


-0.14 


Platinum- 1.10 


Cadmium +0.16 


Hydrogen 


-0.24 


Gold -1.7 



Metals have what is called a solution pressure which tends 
to cause them to pass into solution, that is, to become ionic. 
When a metal high in the series, such as zinc, is placed in the 
solution of a metal lower in the series, such as copper, the 
zinc tends to go into solution as zinc ions, thereby making 
the solution positive. The metallic zinc becomes corre- 
spondingly negative and attracts the positive copper ions 
(already in the solution), which are deposited as metallic 
copper on the zinc. The simplified equation for the whole 
change is 



Zn + Cu 4 



= Zn 4 



+ Cu 



Similar changes occur with other metals and solutions. 

When a metal is placed in a solution of one of its own 
salts, a difference in potential is developed between the 
metal and the solution ; i.e. the metal and solution become 
oppositely charged. If the metal, like zinc, is high in the 
series, the solution is positive, but if well down on the list, 
e.g. copper, the solution is negative. The values in the 
above table are the differences in potential (in volts) between 
the various metals and normal solutions of their ions. These 
values may be used to calculate the electromotive force of a 
cell. Thus, in a gravity cell, which consists essentially of 
a copper electrode in copper sulphate solution and a zinc 
electrode in zinc sulphate solution, the electromotive force 



452 INORGANIC CHEMISTRY 

(for normal solutions) is the difference between the electrode 
potentials, or + 0.52 — (— 0.58) = 1.10. The solutions in 
the ordinary cell are not normal, but the calculated and 
actual values of the electromotive force are nearly the same. 
Hydrogen is not a metal in the common acceptance of 
this term, but hydrogen ions (H + ) are positive, so hydrogen 
is usually included in the electrochemical series of the metals. 
Its position is interesting. Metals that precede hydrogen 
displace it from most acids, while those that follow do so 
rarely, if ever. That is, hydrogen is displaced from its solu- 
tions only by metals having a greater solution pressure. 

Silver 

Silver is one of the oldest metals. For ages it has been 
used in the form of ornaments, costly vessels, and coins. 
It is a noble metal, i.e. one which does not oxidize readily 
in the air. The Latin name of silver is argentum, from 
which the symbol Ag is derived. The alchemists called it 
luna on account of its silvery or "moonlike" appearance, 
and its alchemistic symbol was a crescent. 

Occurrence of Silver.- — Native silver is found in Arizona, 
Mexico, Norway; also in South America and Australia. 
The chief ores are the sulphides. The simple sulphide 
(silver glance or argentite, Ag 2 S) is the richest ore, and is 
found alone in many localities in the United States; it also 
occurs mingled with sulphides of lead, copper, antimony, 
or arsenic. These complex sulphides are found in Mexico, 
Peru, Bolivia, Chile, and in Idaho. Small quantities of 
native silver chloride (horn silver, AgCl) are also found; 
it resembles wax or horn in softness and color. Alloys oi 
silver with gold, mercury, and copper are found; average 
California gold contains about 12 per cent silver. Many 
ores contain silver, especially those of lead; and this argen- 



SILVER 453 

tiferous (or silver-bearing) lead is one of the chief sources of 
silver. 

Metallurgy of Silver. — Silver is extracted from its ores 
by two principal processes. (1) In the amalgamation process 
the powdered ore is first changed into silver chloride by 
roasting (or simply mixing) it with sodium chloride. The 
mass is then reduced to silver by agitation with water and 
iron (or an iron compound) ; the simplest equation for this 
reaction is : — 

2AgCl + Fe = 2Ag + FeCl 2 

Silver Iron Silver Iron 

Chloride Chloride 

The silver is removed by adding mercury, w T hich forms an 
amalgam with the silver, but not with the other substances. 
When the amalgam is heated, the mercury distills off, and 
the silver — with some gold — remains behind. (2) Silver 
is extracted from lead ores by the Parkes process. After 
the sulphur, arsenic, and other impurities have been re- 
moved from the lead ores, the final product is a mixture of 
lead, silver, and gold. This is melted and thoroughly 
mixed with zinc. As the mixture cools, an alloy of silver, 
gold, zinc, and a little lead rises to the top, solidifies, and is 
removed. The remaining lead mixture is treated again 
with zinc. The alloy of silver, gold, zinc, and lead is heated 
to volatilize the zinc and to oxidize (or to melt away) the 
lead. The silver and gold are separated by electrolysis or 
by heating the mixture with sulphuric acid; the gold is not 
acted upon in the latter process, but the silver forms silver 
sulphate, which is reduced by copper to metallic silver. 

Lead ores containing considerable silver are sometimes 
subjected to cupellation to extract the silver. The ore or 
alloy is heated in a furnace having a shallow hearth made 
of porous, infusible bone ash. The lead is thereby changed 
into an oxide (PbO ; litharge), which melts, and is partly 



454 INORGANIC CHEMISTRY 

driven off by the air blast into pots and partly absorbed by 
the porous cupel. The silver is prevented from oxidizing 
by the melted litharge, but toward the end of the operation 
the thin film of litharge bursts, and the metallic silver ap- 
pears as a bright disk, if the operation is conducted in a 
furnace, and as a globule or button, if the extraction is 
performed in a small assay cupel. The process is then 
stopped and the silver removed. 

Properties of Silver. — Silver is a lustrous, white metal, 
which takes a brilliant polish. It is harder than gold, but 
softer than copper. Like copper, it is tenacious, ductile, 
and malleable, and can be easily made into various shapes. 
Its specific gravity is about 10.5. It melts at about 962° C, 
and fuses readily on charcoal in the blowpipe flame; it 
vaporizes in the oxy hydrogen flame and in the electric fur- 
nace. Molten silver absorbs about twenty times its volume 
of oxygen, which is expelled violently when the silver solidifies. 
Silver is an excellent conductor of heat and electricity, but 
it is too expensive for such uses. It does not tarnish in 
air, unless sulphur compounds are present (especially hydro- 
gen sulphide), and then the familiar brown or black film of 
silver sulphide is produced. This blackening is especially 
noticed on silver spoons which have been put into eggs or 
mustard, and on silver coins which have been carried in 
the pocket, the sulphur in the latter case coming from the 
sulphur compounds in the perspiration; the tarnishing of 
household silver is due to sulphur compounds in illuminat- 
ing gas or the gas from burning coal. Tarnished silver can 
be quickly cleaned by placing it in contact with a piece of 
aluminium in a hot solution of sodium bicarbonate and 
sodium chloride. " Oxidized " silver is not oxidized, but 
coated with silver sulphide. Silver is only very slightly 
acted upon by hydrochloric acid, and not at all by melted 



SILVER 



455 



potassium hydroxide, sodium hydroxide, or potassium 
nitrate. Nitric acid and hot concentrated sulphuric acid 
change it into the nitrate and sulphate respectively. 

Alloys of Silver. — Pure silver is too soft for constant 
use, so is usually hardened by adding a small amount of 
copper. These alloys are used in making coins and jewelry. 
The silver coins of the United States and France contain 
900 parts of silver to 100 of copper, and are called 900 fine. 
British silver coins are 925 fine; this quality is called " ster- 
ling silver/ ' and from it much ornamental and useful silver- 
ware is made. 



IC" 



Silver Plating. — Metals cheaper than silver can be coated 
or plated with pure silver in largely the same way as copper. 
Plated silverware has the ap- 
pearance of solid or pure sil- 
ver. The object to be plated 
is carefully cleaned, and made 
the cathode in an electrolytic 
cell containing a solution of 
potassium silver cyanide 
(KAg(CN) 2 ). The anode is 
a plate of pure silver (Fig. 70). Silver-coated mirrors are 
made by reducing ammonio-silver nitrate with an organic 
compound, e.g. formaldehyde. 




Fig. 70. — Apparatus for silver plat- 
ing. A, A, A, are silver anodes, 
and the spoons are cathodes. 



Compounds of Silver. — The most important compound 
is silver nitrate (AgN0 3 ). It is a white, crystalline solid, 
readily soluble in water. It is made by treating silver with 
nitric acid. The equation for the chemical change is : — 



3 Ag + 4 HN0 3 

Silver Nitric 

Acid 



3 AgN0 3 + NO 

Silver Nitric 

Nitrate Oxide 



+ 2 H 2 

Water 



456 INORGANIC CHEMISTRY 

Exposed to the light, it turns dark if in contact with organic 
matter. It discolors the skin; if applied long enough, it 
disintegrates the flesh, and is often used by physicians for 
this purpose. Its caustic action and the silvery color of 
the metal from which it is made long ago led to the name, 
lunar caustic. Besides its extensive use in photography and 
silver plating, silver nitrate is the essential constituent of 
some indelible inks. Silver chloride (AgCl) is made by add- 
ing hydrochloric acid or any soluble chloride to a solution 
of a silver compound. Thus formed, it is a white, curdy 
solid, which turns violet in the light, and finally black. 
This action of light is more intense if organic matter is 
present. Silver chloride is only slightly soluble in water, 
but it dissolves in ammonium hydroxide, in sodium thio- 
sulphate solution, and in potassium (or sodium) cyanide 
solution; in each case a complex compound is produced. 
'The formation and properties of silver chloride constitute 
the test for silver. Silver bromide (AgBr) and silver iodide 
(Agl) are similar to silver chloride in their properties. Silver 
bromide is slightly yellow and silver iodide has a distinct 
yellow tinge; compared with the chloride, both are less 
soluble in ammonium hydroxide; the bromide dissolves 
readily and the iodide only slightly in sodium thiosulphate 
solution. Silver bromide and iodide find extensive applica- 
tion in photography. 

Silver compounds, if soluble in water, yield silver ions 
(Ag + ). Complex ions are formed with other solvents, e.g. 
Ag(NH 3 ) 2 + with an excess of ammonium hydroxide and 
Ag(CN) 2 " with an excess of potassium cyanide solution. 

The atomic weight of silver is 107.88 and the valence is one. 

Photography is based on the fact that silver salts, es- 
pecially the bromide, change chemically when mixed with 
organic matter and exposed to the light. The photograph 



GOLD 457 

is taken on a glass plate, coated on one side with a thin 
layer of gelatine, containing the silver salt. Sometimes a 
sheet of sensitized gelatine, called a film, is used. The plate 
or film is placed in the camera and exposed. The light, 
which is reflected from the object being photographed, 
changes the silver salt in proportion to its brilliancy. The 
plate, however, shows no change until it has been developed. 
This process consists in treating the plate w T ith a reducing 
agent, e.g. hydroquinone, pyrogallic acid, or special mix- 
tures. As the developer acts upon the silver salt, the image 
appears. This is really a deposit of finely divided silver. 
Where the intense light fell upon the plate, the deposit is 
heavier than where little or no light fell. Hence, dark 
parts of the object appear light on the plate, and light parts 
dark ; and since the image is the reverse of the object, the 
plate is called a negative. When the plate has been properly 
developed, it still contains some silver salt not altered by 
the light ; and if it were left on the plate, the image would 
be clouded and finally obliterated by the light. The image 
is, therefore, fixed by washing off the unreduced silver salt 
with a solution of sodium thiosulphate (or "hyposulphite"). 
A print is made by laying sensitized paper upon the nega- 
tive and exposing them to the sunlight, so that the light 
will pass through the negative to the paper. The negative 
obstructs the light in proportion to the thickness of the 
silver deposit, so the photograph has the same shading as 
the object. Most prints, like the plates, must be fixed. 
Sometimes the color is improved by toning, i.e. by placing 
the print in a solution of gold or of platinum. 

Gold 

Gold, like silver, is one of the oldest metals. For ages it 
has been the most highly prized of the metals and exten- 



458 INORGANIC CHEMISTRY 

sively used for personal adornment and for the fundamental 
standard of value. Chemically it is a noble metal. 

The Latin name of gold, aurum, gives the symbol Au. 
For several centuries the mediaeval chemists or alchemists 
tried to produce gold by the transmutation of base or cheaper 
metals. They were unsuccessful in their search for the 
philosopher's stone, which they believed had power to effect 
this transformation. 

Occurrence of Gold. — Gold is widely distributed in the 
native state, but not abundantly in many places. Unlike 
copper and silver, its compounds are few and rare; the 
only important ones are the tellurides (compounds of tel- 
lurium, e.g. AuAgTe 2 ) found in Colorado. Gold is never 
found pure, being alloyed with silver and occasionally with 
copper or iron. It is disseminated in fine, almost invisible, 
particles among ores of other metals, especially the sul- 
phides of copper and iron, though not so abundantly as 
silver. Much gold is found in veins of quartz, and in the 
sand and gravel formed from gold-bearing rocks. Gold 
occurs usually as dust, scales, or grains, but occasionally 
shapeless masses called "nuggets"-" are' found, varying in 
weight from a few grams to many kilograms. The largest 
nugget ever known weighed about 84 kg. (184 lb.). 

The chief gold-producing countries are the United States, Aus- 
tralia, South Africa, and Russia. The annual production in the 
United States is about four million troy ounces, which comes largely 
from Colorado, California, other Western states, and Alaska. 

Gold Mining and Metallurgy. — Gold was first obtained 
by miners by washing the gold-bearing sand and gravel of 
a stream in large pans or cradles. This primitive method 
was soon replaced by placer mining and hydraulic mining. 
Streams of water, directed against the earth containing the 



GOLD 459 

gold, wash away the lighter materials, but leave the heavy 
gold behind as fine particles called gold dust. From this 
mixture gold and silver are extracted by mixing with mer- 
cury, or by passing the moistened mass over copper plates 
coated with mercury. The amalgam is then heated, as 
in the metallurgy of silver, to remove the mercury; the 
residue of gold and silver is separated as described below. 
In vein mining the gold-bearing rock — usually quartz — 
is crushed and then washed, and the gold removed by mer- 
cury, as in placer mining. Low grade ores and those con- 
taining certain metals cannot be profitably treated with 
mercury. In the chlorination process the crushed ore is 
roasted and then revolved in barrels containing bleaching 
powder and sulphuric acid; this operation forms a soluble 
gold chloride (AuCl 3 ), from which the gold is precipitated 
as a fine powder by ferrous sulphate (or other reducing 
agents). Sometimes liquid chlorine is used in the chlorina- 
tion process. In the cyanide process the crushed ore, or the 
slime from a previous extraction, is mixed with a weak solu- 
tion of potassium (or sodium) cyanide in large vats exposed 
to the air; this operation changes the gold into a soluble 
cyanide, thus : — 

4 Au + 8 KCN + 2 + 2 H 2 = 4 KAu(CX) 2 + 4 KOH 

Gold Potassium Oxygen Water Potassium Potassium 

Cyanide Gold Cyanide Hydroxide 

The gold is precipitated as a purple powder from this solu- 
tion by electrolysis or by treatment with zinc. 

Purification of Gold. — Gold obtained by the foregoing 
methods is impure, — silver, copper, and lead being the chief 
impurities. The purification of gold is accomplished by 
chemical or electrolytic processes. By the old parting 
process known as quartation an alloy of gold and silver, in 
which gold is about one fourth of the whole, is treated with 



460 INORGANIC CHEMISTRY 

nitric acid; this operation changes the silver into the nitrate, 
from which the pure gold can be readily removed. The 
metals can also be parted by the cheaper method described 
under silver, viz. by boiling with concentrated sulphuric 
acid. By this treatment the gold, which is about one sixth 
of the alloy, is left as a brownish, porous mass. It is washed 
and dried, and then fused into a coherent mass with char- 
coal and sodium carbonate. These chemical processes have 
been largely displaced by electrolytic methods of separation. 
In one of the latter the anode is an alloy of gold and silver 
(low in gold), the cathode is silver, and the electrolytic solu- 
tion is a dilute silver nitrate solution containing a small pro- 
portion of nitric acid. When the current passes, part of the 
silver of the anode goes into solution as the nitrate, while part 
is deposited at the cathode; the gold remains at the anode as 
a fine powder and is caught in a cloth bag which incloses the 
whole anode. In another electrolytic process, which is 
successfully operated in the United States mints, the anode 
is an alloy rich in gold, the cathode is pure gold, and the 
electrolytic solution is a solution of gold chloride (AuCl 3 ) 
and hydrochloric acid; very pure gold is deposited on the 
cathode, while the silver forms silver chloride around the 
anode. 

The purity of gold is expressed in carats. Pure gold is 24 
carats fine; an alloy containing 22 parts of gold and 2 parts 
copper is 22 carat gold, while one containing equal parts gold 
and other metals is 12 carat gold. 

Properties of Gold. — Gold is a lustrous, yellow metal. 
It is about as soft as lead, and is the most ductile and malle- 
able of all metals. The leaf into which it is readily beaten 
is very thin. The melting point is 1063° C. Gold is one of 
the heaviest metals, its specific gravity being about 19. Air, 
oxygen, hydrogen sulphide, water vapor, and the common 



GOLD 461 

acids do not attack it; but it is changed into a gold chloride 
(AuClo) by chlorine water and compounds and mixtures 
which liberate chlorine, especially a mixture of concentrated 
nitric and hydrochloric acids. This mixture, long known 
as aqua regia, has already been discussed. (See Aqua Regia.) 

Uses of Gold. — Pure gold is too soft for most practical 
uses, and is therefore usually hardened by adding copper or 
silver. The gold-copper alloy has a reddish color and is 
often called " red gold "; the gold-silver alloy is paler than 
pure gold and is sometimes called " white gold." Gold coins 
contain gold and copper. The United States standard gold 
coins contain 9 parts gold and 1 part copper, while in England 
the legal standard is 11 of gold to 1 of copper. Gold leaf of 
various grades is used to ornament books, signs, and other 
objects. Jewelers use gold for many purposes; such gold 
varies from 12 to 22 carats in purity. On account of its 
malleability, feeble chemical action, and beauty, gold is used 
by dentists for filling teeth. 

Compounds of Gold are readily decomposed by metals, 
reducing agents {e.g. ferrous sulphate or hydrogen sulphide), 
fine solids like charcoal, and by electrolysis. When gold is 
dissolved in aqua regia and the excess of acid removed by 
evaporation, the resulting auric chloride (AuCl 3 ) gives with 
stannous chloride solution a beautiful purple precipitate, 
which has long been called " purple of Cassius " ; it is col- 
loidal gold. Its formation is a test for gold. The process 
of gold plating is the same as silver plating, only the solution 
is one of potassium auricyanide (KAu(CN) 4 ) and the anode 
is gold. Much cheap jewelry is gold plated. 

The atomic weight of gold is 197.2. 

The Copper Family belongs to Group I in the periodic 
classification, but it is not a typical family. Not only are 



462 INORGANIC CHEMISTRY 

these elements quite different from the other family in this 
group (i.e. the alkali metals), but the different members are 
not closely related to each other. Indeed copper more closely 
resembles mercury and iron than it does silver, while gold 
has properties which strongly suggest aluminium and iron. 
The valence of copper, as already stated, is one in cuprous 
compounds and two in cupric. The valence of silver is 
usually one. Gold has the valence of one in aurous com- 
pounds and three in auric compounds. 

Problems and Exercises 

1. How much silver chloride is formed by adding 10 gm. of 
crystallized barium chloride (BaCl 2 . 2 H 2 0) to silver nitrate? 

2. How many pounds of copper can be obtained from a ton 
(2000 lb.) of pure chalcopyrite (CuFeS 2 ) ? 

3. A flask filled with water weighed 153 gm. ; 25 gm. of cop- 
per were dropped in. The flask and contents then weighed 175.19 
gm. What is the sp. gr. of copper? 

4. What weight of gold will 25 gm. of ferrous sulphate precipi- 
tate from a solution of auric chloride? (Equation is AuCl 3 + 
3 FeS0 4 = Au + Fe 2 (S0 4 ) 3 + FeCl 8 .) 

5. What weight of (a) silver and (h) gold will be precipitated 
from the respective solutions by 25 gm. of copper? 

6. Starting with silver, how would you prepare in succession 
silver nitrate, silver chloride, silver? 

7. As in Exercise 6, the f olio wing from copper : the oxide, 
nitrate, oxide, metal, sulphate, metal? 

8. Calculate the atomic weight of copper, silver, or gold from 
the following : (a) 20.6885 gm. of copper oxide give 16.5167 gm. 
of copper; (6) 4.39313 gm. of copper precipitated 14.9104 gm. of 
silver ; (c) the specific heat of gold is .032. 

9. Calculate the solubility products of silver chloride, bromide, 
and iodide if the ionization is 100 per cent and the molar solubilities 
are .0000094, .00000058, and .000000016 respectively. 

10. Write the formulas of the following compounds by apply- 
ing the principle of valence (see Chapter XIV) : Cuprous iodide, 
cupric acetate, silver cyanide, silver sulphate, silver oxide, aurous 
bromide, auric chloride, aurous hydroxide, cupric hydroxide. 



CHAPTER XXVII 

Calcium, Strontium, and Barium — Radium 

These elements form a natural family in Group II of the 
periodic classification known as the alkaline earth metals. 



Calcium 

Occurrence. — Calcium is never found free. Combined 
calcium makes up about 3.5 per cent of the earth's crust. 
The most abundant compound is cal- 
cium carbonate (CaC0 3 ) . Many rocks 
are silicates of calcium and other 
metals. Calcium sulphate (CaS0 4 ) 
is abundant. Calcium compounds 
are essential to the life of plants 
and animals, being found in the 
leaves of plants, and in the bones, 
teeth, and shells of animals. The 
ocean, many rivers, lakes, and springs 
contain calcium salts, especially the 
acid carbonate (H 2 Ca(C0 3 )2) and the 
sulphate. 




Fig. 70 a. — Apparatus 
for preparing calcium. 



Preparation. — Calcium is prepared 
by the electrolysis of melted calcium 
chloride (Fig. 70 a). The anode is a 

graphite crucible (A) and the cathode is a rod of iron 
(JS) which can be adjusted by a screw (C) so that it will 

463 



464 INORGANIC CHEMISTRY 

just dip into the melted chloride. The lower part of the 
crucible is cooled by a current of water in EE. When the 
current passes calcium collects on the end of the cathode 
and is gradually removed as an irregular rod (D) by slowly 
raising the cathode ; the end of the calcium rod dips into 
the electrolyte and thus serves as the lower end of the 
original cathode. 

Properties. — Calcium is a silvery metal. Its specific 
gravity is about 1.55 and its melting point is about 810° C. 
It tarnishes in the air and should be kept under some water- 
free liquid. When heated in air it forms an oxide (CaO) 
and a nitride (Ca 3 N 2 ) . It interacts with water, slowly at 
ordinary temperatures and rapidly at high temperatures, 
forming calcium hydroxide (Ca(OH) 2 ) and hydrogen. With 
acids it interacts readily, yielding hydrogen and a calcium 
salt. 

Calcium Carbonate, CaC0 3 . — The most abundant form 
of this compound is limestone. Vast deposits are found in 
many places. Limestone is a white or gray compact solid, 
but impurities, especially organic matter and iron compounds, 
produce blue, yellow, reddish, and black varieties. Hard, 
crystalline limestone which takes a good polish is called 
marble. This form, which has a wide range of color, is used 
as a building and an ornamental stone. Calcite is crystal- 
lized calcium carbonate. It is almost as abundant as quartz, 
though softer; its varied color and crystal form combine to 
make it an attractive mineral (Fig. 71). A very transparent 
variety of calcite called Iceland spar has the remarkable 
property of double refraction, i.e. making small objects 
appear double. 

Calcium carbonate is not soluble in water, unless carbon 
dioxide is present. (See also Carbon Dioxide and Acid Cal- 



CALCIUM 465 

cium Carbonate.) As water containing carbon dioxide works 
its way underground in limestone regions, the limestone is 
dissolved and caves are often formed or enlarged. When 
the water enters a cave and drips from the top, the water 
evaporates, or the gas escapes, or both, and the calcium car- 
bonate is redeposited, often forming stalactites and stalag- 
mites. The stalactites hang from the roof like icicles,_while 





Fig. 71. — Calcite crystals. 

the stalagmites grow upon the floor, as the deposit slowly 
accumulates from the solution which drops from the roof 
or the tips of stalactites. The Mammoth Cave in Kentucky, 
the Marengo Cave in Indiana, and the Luray Cavern in Vir- 
ginia are famous for these fantastic formations. Mexican 
onyx is a variety of stalagmite. Vast deposits of this beauti- 
ful mineral are found in Algeria and Mexico. It is translu- 
cent and delicately colored, and is used as an ornamental 
stone, especially for altars, table tops, mantels, soda foun- 
tains, and lamp standards. Deposits of limestone are found 
around many mineral springs. Travertine occurs near many 
springs in Italy. When fresh, it is soft and porous, but it 
soon hardens and becomes a durable building stone in dry 
climates. The w T alls of many ancient buildings in Italy 
are travertine. Limestone often contains shells and fossils, 
confirming our belief that limestone is the remains largely 
of the shells of animals. The calcium carbonate dissolved 



466 



INORGANIC CHEMISTRY 



in the ocean is transformed by marine organisms into shells 
and bony skeletons. The hard parts of these animals accu- 
mulate in vast quantities on the ocean bottom, become com- 
pact, and are finally elevated above the surface of the ocean 
On certain parts of the coast of Florida, coquina or shell 
rock is found. It is a mass of fragments of shells cemented 
by calcium carbonate, and in time it may become compact 





Fig. 72. — Ooze from the ocean bot- Fig. 73. — Globigerina shells (mag- 
tom, showing globigerina shells nified) found in chalk from Iowa, 

(magnified). 

limestone. Chalk is the remains of shells of minute animals. 
When examined under the microscope, a good specimen is 
seen to consist almost entirely of tiny shells. The ocean 
contains myriads of minute animals, and when they die, 
their shells, which are calcium carbonate, sink to the bottom. 
As a result, the ocean bottom is partly covered with a gray 
mud, called globigerina ooze. Under the microscope this 
ooze has the appearance shown in Figure 72, and when dried 
and compressed it can hardly be distinguished from chalk. 



CALCIUM 467 

It is believed that the immense beds of chalk found in Eng- 
land and other places were formed from this ooze. Some 
varieties of chalk, under the microscope closely resemble the 
ooze (Fig. 73). Whiting is a variety of impure chalk; putty 
is a mixture of whiting and oil. Coral is calcium carbonate, 
and the vast accumulations in the sea are the skeletons 
of the coral animals. 

The properties of calcium carbonate, discussed in Chapter 
XV, may be profitably reviewed by the student at this point. 

Calcium Bicarbonate or Acid Calcium Carbonate, Ca(HC0 3 ) 2 
or H 2 Ca(C0 3 ) 2 , is a salt formed by dissolving calcium car- 
bonate in water containing carbon dioxide. It decomposes 
readily and deposits the normal carbonate (CaC0 3 ). (See 
also Stalactites and Hardness of Water.) 

Calcium Oxide, CaO, is the chemical name of lime. It is a 
porous white solid. Pure lime is almost infusible, and when 
heated in the oxyhydrogen flame it gives an intensely bright 
light, sometimes called the "lime light. " (See Hydrogen and 
Calcium Light.) In the electric furnace it melts and vol- 
atilizes, if the heating is prolonged. Lime containing im- 
purities, like sand, clay, and iron compounds, melts quite 
readily into a glass or slag. Exposed to the air, lime becomes 
" air slaked," that is, it slowly absorbs water and carbon 
dioxide, swells, and soon crumbles to a powder, which is a 
mixture of calcium hydroxide and calcium carbonate. When 
fresh lime and water are mixed, the lime soon becomes 
warm, swells, and finally liberates considerable heat, as may 
often be seen when mortar is being prepared. This opera- 
tion is called " slaking," and the product is " slaked lime." 
The equation for the chemical change is — 

CaO + H 2 = Ca(OH> 

Calcium Water Calcium 

Oxide Hydroxide 



468 



INORGANIC CHEMISTRY 



Fresh lime disintegrates organic matter, and is therefore 
often called "caustic lime" or quicklime. It combines with 
water to form calcium hydroxide and interacts with acids 
to form calcium salts. 

Lime is one of the most useful substances. It is used in 
preparing mortar, cement, and metals, in making bleaching 
powder, calcium carbide, and glass, in purifying illuminating 
gas and sugar, in removing hair from hides before the process 
of tanning, in dyeing and bleaching cotton cloth, in drying 
gases, and as a disinfectant and fertilizer. 

Lime is prepared on a large scale by heating limestone in 
a partly closed cavity or vessel. The decomposition takes 
place according to the equation — 

CaC0 3 = CaO + CO, 

Carbon 



Calcium 
Carbonate 



Calcium 
Oxide 



Dioxide 



The carbon dioxide gas escapes, and the lime is left in the 

kiln. 

Limestone was formerly 
"burned" in a cavity on a 
hillside, and in some regions 
it is so prepared to-day. This 
is a crude process and has 
been largely superseded by 
well-regulated decomposition 
in limekilns. The older or 
periodic kiln is constructed of 
bricks or of blocks of lime- 
stone and loosely filled with 
pieces of limestone from the 
arch to the top, as shown in 
Figure 74. The fire is built 




Fig. 74. 



■ Periodic limekiln (vertical 
section). 



at the bottom and burns several days ; when the kiln is cool, 
the lime is removed. These kilns have been largely replaced 



CALCIUM 



469 



by continuous kilns (Fig. 75). Limestone is introduced at 
A and decomposed by the heat from the gases generated in 
B y B. The lime is withdrawn at C, C. Carbon dioxide 
and other waste gases escape 
through the top of the kiln. 

Cement is made by heating 
a pulverized, carefully propor- 
tioned mixture of limestone, 
clay, and sand. This mixture 
is fed into the upper end of a 
long, inclined, tubular furnace, 
which is heated by powdered 
coal blown in at the lower 
end ; as the furnace revolves 
slowly the mixture interacts, 
moves along, and finally drops 
out as "clinker," which con- 
sists essentially of calcium sili- 
cate, calcium aluminate, and an excess of calcium oxide. 
The " clinker " is pulverized and mixed with ground gypsum. 
Cement and water slowly interact and the mixture sets to a 
hard mass ; it sets under water as well as in air. Concrete 
is a mixture of cement, water, sand, and crushed stone. 
Cement and concrete are used as construction material. 

Soda Lime is a mixture of lime and sodium hydroxide. 

Calcium Hydroxide, Ca(OH) 2 , is a white powder. It is 
sparingly soluble in water, but more soluble in cold than in 
warm water. The solution has a bitter taste, an alkaline 
reaction, and is commonly called limewater. Exposed to 
the air, limewater becomes covered with a thin film of cal- 
cium carbonate, owing to the interaction of carbon dioxide 
and calcium hydroxide. For the same reason, limewater 




Fig. 75. — Continuous limekiln. 



470 INORGANIC CHEMISTRY 

becomes turbid or cloudy when carbon dioxide is passed into 
it. The formation of calcium carbonate in this way is the 
usual test for carbon dioxide. The equation for this chemical 
change is — 

H 2 



Ca(OH) 2 


+ co 2 = 


CaC0 3 + 


Calcium 


Carbon 


Calcium 


Hydroxide 


Dioxide 


Carbonate 



Limewater is prepared by carefully adding lime to consider- 
able water, allowing the mixture to stand for a day or two 
in a stoppered bottle, and then removing the clear liquid. 
When considerable calcium hydroxide is suspended in the 
liquid, the mixture is called milk of lime. Ordinary white- 
wash is thin milk of lime. Limewater is used in the chemical 
laboratory and as a medicine. 

Mortar is a thick paste formed by mixing lime, sand, and 
water. It slowly hardens or "sets," owing to the loss of 
water and the absorption of carbon dioxide. It hardens 
without much shrinking, and when placed between bricks 
or stones holds them firmly in place. The sand makes the 
mass porous and thus facilitates the change of the hydroxide 
into the carbonate. The sand is changed chemically only 
to a very slight extent, if at all. Hair is sometimes added to 
make the mortar stick better, especially when it is used as 
plaster for walls. 

Calcium Sulphate, CaS0 4 . — This salt occurs abundantly 
and in several varieties. Anhydrite (CaS0 4 ) is associated 
with sodium chloride. The other varieties are often grouped 
under the general term gypsum and have the composition 
CaS0 4 . 2 H 2 0. Ordinary gypsum occurs as white masses 
and is known as rock or massive gypsum. The lustrous, 
translucent, crystalline kind is called selenite and is very 
pure; fine-grained masses are named alabaster, and the 



CALCIUM 471 

fibrous varieties are satin spar. Most varieties of gypsum 
are soft, selenite being easily scratched with the finger nail. 

Calcium sulphate is slightly soluble in water. When gyp- 
sum is heated, it gradually loses its water of crystallization, 
and becomes opaque (if previously crystalline) and friable. 
When this dehydrated product is mixed with water, it forms 
a paste which soon solidifies or "sets" to a white porous 
mass with a smooth surface. If the raw material is of good 
quality and the temperature is kept near 125° C, the final 
product is plaster of Paris; it has the composition 
(CaS0 4 ) 2 .H 2 and sets quickly owing to the formation of a 
network of crystals of the less soluble salt (CaS0 4 . 2 H 2 0). 
But if the gypsum is impure or the temperature high, the 
product sets slowly; this variety is called cement plaster. 
Sometimes a slowly setting plaster is made by adding a re- 
tarder such as alum or borax during or after the calcination. 
Many kinds of plaster are made from gypsum and their names 
are applied rather indefinitely. 

Plaster of Paris expands slightly in setting and is therefore 
used to make exact reproduction of statues and ornaments; 
considerable is used in making molds and in cementing glass 
to metal. Alabaster, being soft and beautiful in texture, 
is carved into statues and ornaments. Many grades of cal- 
cined gypsum are used as a coating for walls and as the 
principal ingredient of fine plasters. Crude gypsum is used 
in the manufacture of glass and porcelain; pulverized gypsum 
is used as a fertilizer under the name of land plaster. Stucco 
is a mixture of plaster of Paris and glue. 

Calcium Chloride, CaCl 2 . is a white solid. It absorbs mois- 
ture rapidly and is used to dry many gases, not only in the 
laboratory, but also in such industrial processes as the manu- 
facture of carbon dioxide. The crystallized variety dissolves 
readily in water, and the solution is attended by a marked 



472 INORGANIC CHEMISTRY 

fall of temperature. A mixture of crystallized calcium 
chloride and snow produces a temperature of — 40° C. 

The low freezing point of calcium chloride solutions has 
led to their use in refrigerating plants and in automatic 
sprinkler systems. The liquid left from the interaction of 
calcium carbonate and hydrochloric acid contains calcium 
chloride, which on concentration is deposited in large crystals. 
These readily absorb water, but lose their own water of 
crystallization when heated above 200° C. This anhydrous 
calcium chloride is porous and is the form usually used as a 
drying agent. At a high temperature it melts, and solidifies 
in cooling to a hard mass known as fused calcium chloride. 

Calcium chloride is found in small quantities in some of the Stass- 
furt salts. It is obtained in large quantities as a by-product in the 
manufacture of sodium carbonate (by the Solvay process) and other 
chemicals. 

Compounds of Calcium. — Calcium cyanamide (CaN 2 C) 
is made by passing nitrogen over heated calcium carbide ; 
it is used as a fertilizer because it provides nitrogen in a form 
readily assimilated by growing plants. Calcium sulphide 
(CaS) is formed by heating a mixture of gypsum and car- 
bon; like other sulphides, it stains silver brown. Cal- 
cium oxalate (CaC 2 4 ) is a white solid formed by the in- 
teraction of ammonium oxalate and any soluble calcium 
compound; it is insoluble in acetic acid but soluble in 
hydrochloric acid. Its formation and properties serve as 
a test for calcium compounds. Another test is the yellow- 
red color imparted to the Bunsen flame by calcium com- 
pounds, especially the chloride. The spectrum of calcium 
is characterized by a red and a green line. 

Calcium compounds in aqueous solution yield colorless 
calcium ions (Ca ++ ). The atomic weight of calcium is 
40.07, and its valence is two. 



STRONTIUM AND BARIUM 473 

Calcium Compounds and Hardness of Water. — Calcium 
sulphate is slightly soluble in water, and calcium carbonate, 
as we have already seen, is changed into the soluble, unstable 
acid carbonate by water containing carbon dioxide. Water 
having these salts of calcium in solution is called hard water. 
They form sticky, insoluble compounds with soap, and as long 
as water contains such salts, the soap is useless as a cleansing 
agent. Heat decomposes acid calcium carbonate, and the 
hardness due to the acid carbonate is called temporary hard- 
ness, because boiling removes it; temporary hardness can 
also be removed by the addition of some slaked lime. The 
hardness caused by calcium sulphate cannot be removed 
by boiling and is called permanent hardness. Magnesium 
sulphate, like calcium sulphate, also produces permanent 
hardness. Permanent hardness can be removed, however, 
by adding sodium carbonate to the water, because the cal- 
cium sulphate and sodium carbonate interact and produce 
soluble sodium sulphate and insoluble calcium carbonate ; 
the latter can be removed by filtering. On a small scale 
borax or ammonia may be used to remove both kinds of 
hardness. When hard water is used in steam boilers, the 
calcium salts are often deposited as a hard mass known as 
" boiler scale." Soft water, such as rain water, contains 
little or no calcium or magnesium salts. 

Strontium and Barium 

Strontium, Sr, and Barium, Ba, are uncommon metallic 
elements. They resemble calcium closely in their physical 
and chemical properties. The metals themselves never 
occur free, and are hardly more than chemical curiosities. 
Their compounds are abundant, and some are useful. 

Compounds of Strontium. — The native compounds are 
the beautifully crystalline minerals, strontianite (strontium 



474 INORGANIC CHEMISTRY 

carbonate, SrC0 3 ) and celestite (strontium sulphate, SrS0 4 ). 
Strontium oxide (strontia, SrO), like lime, is made by heating 
the carbonate. It unites with water to form strontium 
hydroxide (Sr(OH) 2 ), which is used in the purification of 
beet sugar. Strontium nitrate (Sr(N0 3 ) 2 ) and other salts 
of strontium color a flame crimson, and are widely used in 
making fireworks, especially "red fire." The latter is a 
mixture of potassium chlorate, shellac, and strontium nitrate. 
Several strontium salts are used as medicines. 

The crimson color imparted to the Bunsen flame is the 
usual test for strontium. The spectrum of strontium is char- 
acterized by several red lines and a blue one. 

Compounds of Barium. — The most abundant native 
compounds are witherite (barium carbonate, BaC0 3 ) and 
barite (barium sulphate, heavy spar, barytes, BaS0 4 ). The 
oxides, BaO and Ba0 2 , have already been mentioned as a 
source of oxygen. Barium hydroxide (Ba(OH) 2 ) solution is 
often called baryta water, and it forms insoluble barium car- 
bonate (BaC0 3 ) by interaction with carbon dioxide. Barium 
chloride (BaCl 2 ) is a white crystalline solid. Barium sul- 
phate (BaS0 4 ) is formed as a fine, white, highly insoluble 
precipitate by the combination of barium and sulphate ions. 
It is used in making (and sometimes adulterating) white 
paint. A specially prepared mixture of barium sulphate 
and zinc sulphide, called lithophone, forms the basis of a 
white paint which is superior in some ways to white lead 
paint. Barium sulphate is also used to fill paper and give 
it a gloss. Barium salts color a flame green, and barium 
nitrate (Ba(N0 3 ) 2 ) is extensively used in making fireworks, 
especially " green fire." Commercial barium sulphide (BaS), 
as well as the sulphides of calcium and strontium, shine feebly 
in the dark, after having been exposed to a bright light. On 
account of this property they are used in making luminous 



STRONTIUM AND BARIUM 475 

paint. Barium chromate (BaCr0 4 ) is a yellow solid ob- 
tained by the interaction of a soluble barium compound and 
potassium dichromate. Soluble barium salts, unlike those 
of calcium, are poisonous. 

The green color imparted to the Bunsen flame is the test 
for barium. The spectrum of barium is characterized by 
several green and orange lines. 

Aqueous solutions of strontium and of barium compounds 
contain colorless strontium ions (Sr ++ ) and barium ions 
(Ba ++ ). 

The Alkaline Earth Family. — This is a typical family 
and resembles its contiguous families, the alkali metals and 
the earth metals. The metals themselves, if calcium is taken 
as a type, are less active than sodium and potassium but 
more active than aluminium. Not only have these elements 
properties which are much alike, but they also show a grada- 
tion in properties as the atomic weight increases. All interact 
with cold water, burn in air, and form analogous compounds 
whose properties are strikingly suggestive of family relation- 
ship; on the other hand their densities increase from about 
1.5 for calcium (at. wt. 40.07) through 2.5 for strontium (at. 
wt. 87.63) to 3.57 fcr barium (at. wt. 137.37). All three ele- 
ments form an hydroxide, a carbonate, a nitrate, and a sul- 
phate. The solubility of the chlorides and nitrates is quite 
marked and decreases in the order in which the metals have 
been studied (calcium, strontium, barium); the hydroxides 
are much less soluble and in the opposite order; while the 
sulphates vary widely in solubility, a liter of water dissolving 
about 2 gm. of calcium sulphate and only .0023 gm. of barium 
sulphate. All the compounds of these elements are white, 
except barium chromate, which is yellow. The valence of 
the alkaline earth elements in their compounds is almost 
invariably two. 



476 INORGANIC CHEMISTRY 

Radium 

Radium is a rare element. It is a constituent of certain 
rare uranium-bearing minerals, especially pitchblende and 
carnotite. Pitchblende is found in Bohemia and carnotite 
in Colorado and Utah. 

The proportion of radium in pitchblende and carnotite is 
minute, only a few milligrams to the ton. Nevertheless the 
radium is extracted from these minerals and crystallized as 
radium bromide. The supply of radium ores is very 
limited. 

The general properties of radium compounds show that 
the element belongs to the alkaline earth group. Metallic 
radium, which was first isolated by Madame Curie in 1910, 
closely resembles barium. It is a silvery white, intensely 
active metal. Radium forms compounds analogous to those 
of barium, the best known being the chloride (RaCl 2 ), bro- 
mide (RaBr 2 ), and the sulphate (RaS0 4 ). The bromide is 
the common commercial salt and is the substance usually 
meant by the term radium. Radium compounds color the 
Bunsen flame red, have a different spectrum from other ele- 
ments, are naturally phosphorescent, and produce phos- 
phorescence in various substances ; e.g. barium platiho- 
cyanide (BaPt(CN) 4 ), diamond, willemite, and zinc sulphide. 
Besides the properties just mentioned, radium compounds 
have others which are characteristic and differ from those 
exhibited by most substances. Thus, they spontaneously 
evolve relatively large quantities of heat, affect a photo- 
graphic plate, and discharge an electroscope. These proper- 
ties are called radioactive properties. Similar, though less 
active, properties are possessed by uranium, thorium, and 
other substances ; and radioactivity can be more appro- 
priately discussed under these elements. (See Chapter 
XXXII.) 



RADIUM 477 

Problems and Exercises 

1. If the specific gravity of marble is 2.75, how many grams 
would a cubic meter weigh? 

2. What weight of barium sulphate can be obtained by the 
interaction of barium chloride (BaCl 2 ) and 10 gm. of crystallized 
magnesium sulphate, MgS0 4 . 7 H 2 ? 

3. An analysis of a sample of limestone gave 96.45 per cent 
CaC0 3 , 1.00 per cent Si0 2 , .78 per cent MgO, 1.76 per cent Fe 2 3 
and A1 2 3 . How much pure lime could be made from 100 tons of 
this limestone? 

4. What is. (a) the weight and (b) the volume (at standard con- 
ditions) of the gas liberated by the interaction of water and 10 gm. 
of calcium? 

5. Write the formulas of the following compounds by applying 
the principle of valence (Chapter XIV) : Aeid calcium phosphate, 
calcium nitrate, calcium iodide, calcium bromide, strontium chloride, 
strontium oxide, barium sulphide, barium monoxide, barium chlo- 
rate, calcium phosphate (ortho), strontium sulphide. 

6. Apply Exercise 5 to radium. 

7. Calculate the atomic weight of radium if 2.00988 mg. of 
radium chloride are formed by 2.61099 mg. of radium bromide. 

8. Calculate the solubility product of calcium oxalate if the 
ionization is 96 per cent and the molar solubility is .000044. 

9. As in 8, of calcium sulphate for 52.5 per cent and .015. 



CHAPTER XXVIII 
Magnesium, Zinc, Cadmium, and Mercury 

These elements form a natural family in Group II of the 
periodic classification, though the members are not as 
closely related as in the alkali and alkaline earth families, 

Magnesium 

Occurrence of Magnesium. — Magnesium is never found 
free. In combination it is widely distributed and very 
abundant, constituting about 2.5 per cent of the earth's 
crust. Dolomite is magnesium calcium carbonate 
(CaC0 3 . MgC0 3 or CaMg(C0 3 ) 2 ); it forms whole mountain 
ranges in the Tyrol and vast deposits in many regions. 
Dolomite resembles marble and limestone in its properties. 
Magnesium carbonate (MgC0 3 ) is also abundant. Many of 
the Stassfurt salts contain magnesium, for example, kainite 
(KC1, MgS0 4 . 3 H 2 0), carnallite (KC1, MgCl 2 . 6 H 2 0), and kie- 
serite (MgS0 4 . H 2 0). Magnesium is a component of serpen- 
tine, talc, soapstone, asbestos, meerschaum, and other sili- 
cates. The sulphate (MgS0 4 ) and chloride (MgCl 2 ) are found 
in sea water and in mineral springs. 

Preparation of Magnesium. — Magnesium was formerly 
prepared by reducing the chloride with sodium. It is now 
economically manufactured by the electrolysis of fused 
carnallite. A sketch of the essential parts of the apparatus 
is shown in Figure 76. Carnallite is put into the cylindrical 
iron vessel C, which is the cathode. This is closed by the 

478 



MAGNESIUM 



479 




Fig. 76. — Apparatus for the 
manufacture of magnesium by 
the electrolysis of carnallite. 



air-tight cover through which pass the pipes D, D f for 
conveying inert gases into and out of the apparatus. The 
carbon anode A dips into the 
carnallite and is inclosed by 
the porcelain cylinder B, which 
is provided with a pipe E, for the 
escape of the chlorine liberated 
at the anode. The carnallite is 
kept fused by external heat. 
When the current passes, the 
chlorine liberated at the anode 
escapes through E, and the mag- 
nesium liberated at the cathode 
floats on the fused carnallite and 
is prevented from oxidizing by the inert gas supplied through 
D. The porcelain cylinder B prevents the chlorine from 
escaping into the larger vessel. The molten magnesium is 
carefully removed at intervals. 

Properties of Magnesium. — Magnesium is a lustrous, sil- 
very white metal. It is a light metal, the specific gravity be- 
ing only about 1.75. It is tenacious and ductile, and when 
hot can be drawn into w T ire or pressed into ribbon, the latter 
being a common commercial form. It melts at about 650° C. 
and can be cast into different shapes. At a higher tempera- 
ture it volatilizes. It burns with a dazzling light, producing 
dense white clouds of magnesium oxide (MgO) together with 
some magnesium nitride (Mg 3 N 2 ). It does not tarnish in 
dry air, but in moist air it is soon covered with a film of 
oxide. It liberates hydrogen from acids and from boiling 
water. Heated in nitrogen, it forms magnesium nitride 
(Mg 3 N 2 ). 

Uses of Magnesium. — Magnesium in the form of powder 
is used chiefly in taking flash-light photographs. Small 



480 INORGANIC CHEMISTRY 

quantities are used in making fireworks. The powder and 
ribbon are used in the chemical laboratory as a convenient 
form of the metal. Some is used to reduce rare metals 
from their oxides. 

Magnesium Oxide, MgO, is a white, bulky powder. It is 
formed when magnesium burns in the air, but it is manu- 
factured by gently heating magnesite (MgC0 3 ), just as lime 
is made from limestone. It is often called magnesia, or 
calcined magnesia. The native oxide is the mineral periclase. 
Magnesia dissolves very slightly in water, forming magnesium 
hydroxide (Mg(OH) 2 ). A mixture of magnesia and water, 
with or without magnesium chloride, hardens on exposure 
to the air, and is sometimes used as a cement or artificial 
stone. Native magnesium hydroxide is the mineral brucite. 
Like lime, magnesia withstands a high temperature, and is, 
therefore, used as the lining of electrothermal apparatus, 
metallurgical furnaces, and cement kilns, and as the chief 
ingredient of a protective covering for steam pipes. Mag- 
nesia is used in medicine as a mild alkali and as an antidote 
for poisoning by mineral acids. 

Magnesium Sulphate, MgS0 4 , is a white solid. There are 
several crystalline varieties. The native salt kieserite 
(MgS0 4 .H 2 0), when added to water, changes into Epsom 
salts (MgS0 4 . 7H 2 0). This variety was first found in the 
mineral spring at Epsom, England. It is efflorescent, very 
soluble in water, and its solution has a bitter taste. It is 
used in medicine as a purgative, in manufacturing sulphates 
of sodium and potassium, as a fertilizer in place of gypsum, 
and as a coating for cotton cloth. 

Magnesium Chloride, MgCl 2 , is a white solid. It is a by- 
product in the preparation of potassium chloride. The 
crystallized salt (MgCl 2 . 6 H 2 0) is very deliquescent. Mag- 
nesium chloride hydrolyzes with hot water, forming magne- 



MAGNESIUM 481 

sium hydroxide and hydrochloric acid. Hence water con- 
taining magnesium chloride (e.g. sea water) is not suitable 
for use in boilers. 

Magnesium Carbonate, MgC0 3 , occurs native as magnesite, 
and combined with calcium carbonate as dolomite. Mag- 
nesite is converted by heat into magnesia and carbon diox- 
ide ; the magnesia, as stated above, is very generally used as 
a refractory material, while the gas is utilized in the prepara- 
tion of liquid carbon dioxide or of charged beverages. The 
commercial salt known as magnesia alba, or simply mag- 
nesia, is a basic carbonate (Mg4(OH) 2 (C0 3 )3 • 3 H 2 0). 

It was during an investigation of magnesia alba that Black (1728- 
1799) discovered carbon dioxide and showed the close relation between 
analogous compounds of magnesium and calcium. 

Miscellaneous. — Besides the oxide and sulphate other 
magnesium compounds are used as medicines. Fluid mag- 
nesia is prepared by dissolving magnesium carbonate in 
water containing carbon dioxide or by suspending mag- 
nesium hydroxide in water; it is a mild alkali and laxative. 
Magnesium citrate has a similar action; it is an effervescing 
mixture prepared from sodium bicarbonate, tartaric and 
citric acids, sugar, and magnesium sulphate. When disodium 
phosphate and ammonium hydroxide are added to the solu- 
tion of a magnesium compound, a white crystalline precipi- 
tate of ammonium magnesium phosphate (NH 4 MgP0 4 ) is 
produced; its formation is a test for magnesium. Another 
test (though often indifferent) is the pale pink residue left 
after a magnesium compound has been intensely heated in 
a blowpipe flame and subsequently moistened with cobalt- 
ous nitrate solution. 

Soluble magnesium compounds yield colorless magnesium 
ions (Mg ++ ). The atomic weight of magnesium is 24.32, 
and the valence is two. 



482 INORGANIC CHEMISTRY 

Zinc 

Occurrence of Zinc. — Free zinc is never found. The ores 
of zinc are not numerous, though rather abundant. The 
chief ores are the sulphide (sphalerite or zinc blende, ZnS), 
the carbonate (smithsonite, ZnC0 3 ), the silicate (calamine, 
H 2 ZnjjSi0 5 ), and the red oxide (zincite, ZnO). Franklinite 
(Zn(Fe0 2 ) 2 ) and willemite (Zn 2 Si0 4 ) are complex ores which 
also contain manganese and iron. Gahnite has the com- 
position ZnAl 2 4 . 

Zinc ores are found in Germany, Italy, France, Greece, Austria- 
Hungary, Belgium, England, and the United States. Missouri and 
Kansas contain large deposits of the sulphide, while the other ores 
occur chiefly in New Jersey. 

Metallurgy of Zinc. — Zinc is easily smelted. The ores 
are first roasted to change them into the oxide, thus : — 



ZnCOs 


= 


ZnO 


+ 


co 2 




Zinc 
Carbonate 




Zinc 
Oxide 




Carbon 
Dioxide 




2ZnS + 


3 2 


= 


2 ZnO 


+ 


2S0 2 


Zinc Oxygen 
Sulphide 




Zinc 
Oxide 




Sulphur 
Dioxide 



The oxide is then reduced by heating it with coal. This 
operation is conducted in . fire-clay retorts connected with 
iron or clay receivers into which the vapor passes; at first 
it condenses as a blue-gray powder known as zinc dust, 
which is really a mixture of zinc and zinc oxide; but it finally 
condenses as a liquid, which is drawn off at intervals and cast 
into bars or plates. The impure, massive zinc thus obtained 
is called spelter; it is freed from carbon, lead, iron, cadmium, 
and arsenic by repeated distillation, often under reduced 
pressure. 

Properties of Zinc. — Zinc is a bluish white, lustrous metal. 
Its physical properties vary with the temperature. At oiv 



ZINC 483 

dinary temperatures it is brittle, but at 120°-150° C. it is 
soft and can be rolled into sheets and drawn into wire, while 
its specific gravity rises from 6.9 to 7.2. Zinc which has 
been rolled or drawn does not become brittle upon cooling, 
but remains pliable. At 200° C. it again becomes brittle 
and can be easily pulverized. It melts at about 420° C. 
and boils at about 920° C. Heated in the air above its 
melting point, zinc burns with a bluish green flame, forming 
w^hite zinc oxide (ZnO). Zinc does not tarnish in dry air, 
but in moist air it slowly becomes coated with a hard, co- 
herent film which prevents further action. With hot solu- 
tions of sodium and potassium hydroxides, it forms zincates 
and liberates hydrogen, thus : — 



2KOH + 


Zn 


= H 2 + 


K 2 Zn0 2 


Potassium 
Hydroxide 


Zinc 


Hydrogen 


Potassium 
Zincate 



Commercial zinc interacts with dilute acids and liberates 
hydrogen (except from nitric acid). Pure zinc interacts 
wdth acids if in contact with platinum, iron, or copper, or if 
copper sulphate is added to the mixture. Zinc displaces 
many other metals (e.g. lead, tin, copper, and mercury) from 
their solutions; it is strongly electropositive. (See Displace- 
ment of Metals, Chapter XXVI.) 

The vapor density of zinc requires the molecular weight 
66 (approximately). Since the atomic weight is 65.37, a 
molecule of zinc vapor contains only one atom. 

Uses of Zinc. — Zinc in the form of sticks and plates is 
extensively used as the positive electrode in electric batteries. 
Sheet zinc is used as a lining for tanks, and as a protective 
covering which is often placed behind and beneath stoves. 
Iron dipped into zinc becomes coated with zinc and is called 
galvanized iron; it does not rust easily, and is often used in 
place of zinc for roofs, pipes, cornices, and water tanks. 



484 INORGANIC CHEMISTRY 

Telegraph wire is also galvanized. Zinc dust is sometimes 
used in the cyanide process of extracting gold and in many 
chemical experiments in the laboratory. Brass, German 
silver, and other alloys contain zinc. (See Alloys of Copper.) 
Antifriction metals, which are used for bearings, are alloys 
of zinc; Babbitt's metal, for example, contains approxi- 
mately 69 per cent of zinc, 19 of tin, 4 of copper, 3 of anti- 
mony, and 5 of lead. 

Compounds of Zinc. — Native zinc oxide (ZnO) is red, 
owing to the presence of manganese, but the pure oxide 
is white when cold and yellow when hot. It is formed when 
zinc burns, and is manufactured in this way or by heating 
zinc carbonate. It is often called "zinc white" or "Chinese 
white," and is used to make a white paint which is not dis- 
colored by sulphur compounds (especially hydrogen sulphide), 
and is therefore well suited for painting the walls of a chemi- 
cal laboratory. It is also used as an ingredient of ointments 
and cosmetics. Native zinc sulphide (ZnS) is yellow, brown, 
or black on account of impurities, but the pure sulphide is 
white. The latter is formed as a jelly like precipitate when 
hydrogen sulphide is passed into an alkaline solution of a 
zinc salt ; it dissolves in hydrochloric acid, but not in acetic. 
Zinc sulphide is used in making paint (see page 474). Zinc 
sulphate (ZnSO^ is formed by the interaction of zinc and 
dilute sulphuric acid. Large quantities are made by roast- 
ing the sulphide in a limited supply of oxygen and extract- 
ing the sulphate with water. It is a white, crystalline solid 
(ZnS0 4 . 7 H 2 0), which effloresces in the air, and w T hen heated 
to 100° C. loses most of its water of crystallization. The 
crystallized salt is called white vitriol. It is used in dyeing 
and calico printing, as a disinfectant, and as a medicine. It 
is poisonous, but can be safely used externally to relieve in- 
flammation. Zinc chloride (ZnCl 2 ) is a white, deliquescent 



ZINC 485 

solid, prepared by dissolving zinc in hydrochloric acid and 
evaporating the solution until a sample solidifies on cooling. 
It is used in surgery, and also as a constituent of a mixture 
for filling teeth ; large quantities are used to preserve wood, 
especially railroad ties. By hydrolysis it forms hydrochloric 
acid and basic zinc chloride (Zn(OH)Cl). Zinc hydroxide 
(Zn(OH) 2 ) is formed as a dull white, flocculent precipitate 
by the interaction of sodium or potassium hydroxide and the 
solution of a zinc salt, thus : — 

ZnS0 4 + 2NaOH = Zn(OH) 2 + Na 2 S0 4 

Zinc Sodium Zinc Sodium 

Sulphate Hydroxide Hydroxide Sulphate 

An excess of the alkaline hydroxide changes the zinc hydrox- 
ide into a soluble zincate, thus : — 

Zn(OH) 2 + 2NaOH = Na 2 Zn0 2 + 2H 2 

Zinc Sodium Sodium Water 

Hydroxide Hydroxide Zincate 

Zinc salts are poisonous. 

Tests for Zinc. — The formation of the sulphide or the 
hydroxide and a soluble zincate as above described serves 
as the test for zinc. A green incrustation is produced when 
zinc compounds are heated on charcoal and then moistened 
with a cobaltous nitrate (Co(N0 3 ) 2 ) solution. 

Miscellaneous. — Zinc acts as a metal and a non-metal. 
Thus, in compounds like zinc sulphate, zinc chloride, and 
zinc sulphide it is a metal; solutions of the class contain 
colorless zinc ions (Zn ++ ). In the zincates, such as sodium 
zincate (Na 2 Zn0 2 ), the zinc acts as a non-metal. Zinc forms 
complex compounds with ammonium hydroxide, e.g. ammo- 
nio-zinc hydroxide (Zn(NH 3 )4(OH) 2 ). In solutions of com- 
plex zinc compounds, the zinc is often a part of a complex ion. 

The valence of zinc is two. 



486 INORGANIC CHEMISTRY 



Cadmium 



Cadmium, Cd, is an uncommon metal; certain compounds 
are frequently found in zinc and other ores. It also occurs 
as a sulphide (greenockite, CdS). It is white, lustrous, and 
rather soft. Its specific gravity is about 8.6 and its melting 
point is about 320° C. Cadmium is a constituent of certain 
dental amalgams and fusible alloys. (See Bismuth.) Wood's 
metal contains 12 per cent of cadmium. 

The most important compound is cadmium sulphide (CdS). 
This is a bright yellow solid, formed by adding hydrogen 
sulphide to the solution of a cadmium compound. It is 
used as a pigment. Its formation also serves as a test for 
cadmium. Warm dilute sulphuric acid dissolves cadmium 
sulphide owing to the formation of soluble cadmium sulphate, 
but cadmium sulphide does not dissolve in potassium cyanide 
solution; both operations permit the more or less complete 
separation of cadmium and copper compounds. Cadmium 
hydroxide (Cd(OH) 2 ) is a white solid formed upon the addi- 
tion of sodium hydroxide to the solution of a cadmium com- 
pound; it is insoluble in an excess of sodium hydroxide but 
with ammonium hydroxide it forms soluble Cd(NH 3 )4(OH) 2 . 

Cadmium has a valence of two in all its compounds. Its 
vapor density requires the. molecular weight 112 (approxi- 
mately). Since the atomic weight is 112.4, a molecule of 
the vapor contains one atom. 

Aqueous solutions of cadmium compounds contain color- 
less cadmium ions (Cd ++ ). The element also forms com- 
plex ions, e.g. Cd(NH 3 ) 4 ++ and Cd(CN)r~ 

Mercury 

Occurrence of Mercury. — Native mercury is occasionally 
found in minute globules, but the most abundant ore is 
mercuric sulphide (cinnabar, HgS). The ore is mined in 



MERCURY 487 

Spain, Austria, Russia, Italy, and Mexico; in the United 
States large quantities are obtained in California and Texas. 

Mercury has been known for ages as quicksilver. The Latin name, 
hydrargyrum, which gives the symbol Hg. means literally, "water 
silver," emphasizing the fact, so well known, that mercury looks like 
silver and flows like water. 

Preparation of Mercury. — Mercury is prepared by roasting 
cinnabar in an open furnace or closed retort. In the fur- 
nace the sulphide is transformed by the oxygen of the air 
into mercury and sulphur dioxide, thus : — 

HgS + 2 = Hg + S0 2 

Cinnabar Oxygen Mercury Sulphur 

Dioxide 

In the retort lime is mixed with the ore and the equation for 
the reaction is : — 



4 HgS + 


4CaO : 


= 4Hg + 3CaS + CaS0 4 


Mercuric 


Lime 


Mercury Calcium Calcium 


Sulphide 




Sulphide Sulphate 



In each process the mercury is liberated as a vapor and con- 
densed in a series of cooled chambers. The crude mercury 
is freed from soot and dust and collected into large globules 
by stirring and rubbing; it is farther purified by filtering it 
through charcoal or chamois skin. Metallic impurities are 
removed by distillation or by agitation with dilute nitric 
acid. Mercury is sent into commerce in strong iron flasks, 
holding about 75 pounds. 

Properties of Mercury. — Mercury is a bright, silvery metal, 
and is the only one which is liquid at ordinary temperatures. 
It solidifies at about - 38.7° C, and boils at about 357° C. 
It is a heavy metal, the specific gravity being about 13.59. 
It is slightly volatile even at ordinary temperatures, and the 
vapor is poisonous. Mercury does not tarnish in the air, 



488 INORGANIC CHEMISTRY 

unless sulphur compounds are present. At a high tempera- 
ture it combines slowly with oxygen to form the red oxide 
(HgO). Hydrochloric acid and cold sulphuric acid do not 
affect it; hot concentrated sulphuric acid oxidizes it, and 
nitric acid changes it into nitrates. 

The vapor density of mercury requires the molecular weight 
200 (approximately). Since the atomic weight is 200, a 
molecule of the vapor contains only one atom. 

Uses of Mercury. — Mercury is used in making thermom- 
eters, barometers, manometers, and many other kinds of 
scientific apparatus. Its extensive use in extracting gold and 
silver has been mentioned. (See Amalgamation.) Consider- 
able is used in preparing amalgams, medicines, and explo- 
sives {e.g. fulminating mercury, which is used in percussion 
caps and cartridges). It has also come into use recently as 
an electrode in various electrochemical processes. 

Amalgams are alloys of mercury with other metals. They 
are easily prepared by mixing the constituents. Some- 
times the union is violent, as in the preparation of sodium 
amalgam. Amalgamated zinc is usually used in electric 
batteries to prevent unnecessary loss of the zinc. Tin amal- 
gam is sometimes used to coat mirrors. Amalgams of cer- 
tain metals are used as a filling for teeth. Care should be 
taken, while handling mercury, not to let it come in contact 
with gold rings, since gold amalgam is readily formed. Iron 
is the only common metal which does not form an amalgam. 

Compounds of Mercury. — Mercury, like copper, forms two 
classes of compounds — the mercurous and the mercuric. 
The valence of mercury is one in the mercurous compounds 
and two in the mercuric. Mercuric oxide (HgO) is a red 
powder, produced by heating mercury in air or by heating 
a mixture of mercury and mercuric nitrate. As we havQ 



MERCURY 489 

already seen, mercuric oxide is decomposed by heat into 
mercury and oxygen. A yellow variety is produced by the 
interaction of sodium hydroxide and a mercuric salt, thus: — 

2 NaOH + Hg(N0 3 ) 2 = HgO + 2 NaN0 3 + H 2 

Sodium Mercuric Mercuric Sodium Water 

Hydroxide Nitrate Oxide Nitrate 

Mercurous chloride (Hg 2 Cl 2 or HgCl) is a white, tasteless 
pow T der, insoluble in water. It is formed when a chloride 
and mercurous nitrate interact, but it is manufactured usu- 
ally by heating a mixture of mercuric chloride and mercury. 
Under the name of calomel it is extensively used as a medi- 
cine. Mercuric chloride (HgCl 2 ) is a white, crystalline solid, 
soluble in water and in alcohol. It is prepared by heating a 
mixture of mercuric sulphate and common salt. It is a 
violent poison. The best antidote is the white of a raw egg. 
The albumen forms an insoluble mass with the poison, which 
can then be removed from the stomach. The common name 
of mercuric chloride is corrosive sublimate. It has strong 
antiseptic properties and is extensively used in surgery to 
protect wounds from the harmful action of germs; taxider- 
mists sometimes use it to preserve skins, and it has many 
serviceable applications as a medicine and disinfectant. It 
is usually used as a dilute solution (1 part to 1000 parts of 
water). Native mercuric sulphide or cinnabar (HgS) 
is a red crystalline solid. When hydrogen sulphide is passed 
into a solution of a mercuric salt, mercuric sulphide is 
formed as a black precipitate; this variety, when heated, 
changes into red crystals. Vermilion is artificial mercuric 
sulphide. It is manufactured either (1) by grinding together 
mercury and sulphur and treating this mass with potassium 
hydroxide, or (2) by heating mercury and sulphur in iron 
pans and subliming the black mass. In both processes the 
product must be carefully ground, washed, and dried. 



490 



INORGANIC CHEMISTRY 



Chinese vermilion is the best quality. Vermilion has a 
brilliant red color, and, although expensive, is widely used 
to make red paint. Mercurous nitrate (HgN0 3 or Hg 2 (N0 8 ) 2 ) 
and mercuric nitrate (Hg(N0 3 ) 2 ) are prepared by treating 
mercury respectively with cold dilute nitric acid and warm 
concentrated nitric acid. They are white, crystalline solids. 
Soluble mercurous compounds yield mercurous ions (Hg + ) 
and mercuric compounds mercuric ions (Hg ++ ). Both 
kinds are colorless. Mercury also forms complex ions, e.g. 
Hg(CN)r" 



Tests for Mercury. — Clean copper becomes coated with a 
bright film of mercury when put into the solution of any mer- 
cury compound. Other metals act similarly. (See Displace- 
ment of Metals, Chapter XXVI.) Hydrochloric acid precipi- 
tates white mercurous chloride from solutions of mercurous 
salts. Stannous chloride in excess reduces mercuric salts 
at first to white mercurous chloride and finally to a gray 
precipitate of finely divided mercury, thus : — 



2HgCl 2 + 

Mercuric 
Chloride 

2HgCl + 

Mercurous 
Chloride 



SnCl 2 

Stannous 
Chloride 

SnCl 2 

Stannous 
Chloride 



2HgCl 

Mercurous 
Chloride 

2Hg 

Mercury 



+ SnCU 



Stannic 
Chloride 



+ SnCl 4 



Stannic 
Chloride 



This is a typical illustration of the broad use of the terms 
"oxidation" and "reduction," viz. the addition and removal 
of a negative element. In these two chemical changes the 
stannous chloride is oxidized to stannic chloride, i.e. the 
negative element chlorine is added chemically to stannous 
chloride. On the other hand, mercuric and mercurous 
chloride both lose chlorine, becoming respectively mercu- 
rous chloride and mercury, i.e. the negative element chlorine 
is removed chemically from the two mercury compounds. 



ZINC FAMILY 491 

Mercurous salts yield a black precipitate (e.g. ammono-basic 
mercuric chloride and mercury (HgNH 2 Cl + Hg)) with 
ammonium hydroxide, while mercuric salts yield a white one 
(e.g. ammono-basic mercuric chloride (HgNH 2 Cl)). 

The Zinc Family. — The four elements just considered 
(together with beryllium) constitute a natural family in the 
periodic classification. They bear certain resemblances to 
the calcium family, which is in the same periodic group, and 
the different members resemble each other; but the family 
is not a unit, nor does it exhibit the progressive change in 
properties which characterizes certain families. Zinc and 
cadmium are much alike, while mercury differs somewhat 
from these metals and resembles copper. As already stated 
these metals have the valence two, except mercury, whose 
valence is one in mercurous compounds and two in mercuric 
compounds. 

Problems and Exercises 

1. How much magnesium will be formed by heating 100 gm. of 
potassium with magnesium chloride ? (Assume 2 K + MgCl2 = 
Mg + 2KC1.) 

2. What is the per cent of magnesium in (a) magnesite (MgCOs), 
(b) dolomite (MgCa(C0 3 ) 2 ), (c) Epsom salts (MgS0 4 .7H 2 0)? 

3. What is the per cent of zinc in (a) zinc sulphate (ZnSC>4), 
(b) zinc sulphide (ZnS), (c) zinc chloride (ZnCk), (d) zinc oxide (ZnO)? 

4. How much zinc sulphate can be prepared from 65 gm. of zinc? 
from 130 gm. ? from 720 gm. ? 

5. How much mercury is formed by decomposing 400 gm. of cinna- 
bar? (Assume HgS + 2 = Hg 4- S0 2 .) 

6. What is the per cent of mercury in (a) mercuric oxide (HgO), 
(b) calomel (HgCl), (c) corrosive sublimate (HgCl 2 )? 

7. What is the approximate specific heat of cadmium, accepting 
112 as its atomic weight? 

8 If 64.2501 gm. of cadmium sulphate yield 44.4491 gm. of cad- 
mium sulphide, what is the atomic weight of cadmium (assuming 
O = 16andS = 32.07)? 



492 INORGANIC CHEMISTRY 

9. If 380.5744 gm. of mercuric oxide yield 352.4079 gm. of mer- 
cury, what is the atomic weight of this element ? 

10. Write the formulas of the following compounds by applying 
the principle of valence and utilizing analogous formulas in this 
chapter: Magnesium iodide, ammonium magnesium phosphate, 
magnesium fluoride, magnesium silicate, potassium zincate, zinc 
acetate, zinc iodide, cadmium iodide, cadmium sulphate, mercurous 
bromide, mercuric iodide, mercuric sulphate. 

11. What volume of mercury is needed to fill one of the commer- 
cial flasks holding 75 lb. ? 

12. The freezing point of mercury is — 39.5° C. and the boiling 
point is 357° C. What are the corresponding Fahrenheit tempera- 
tures ? 

13. Give the name and formula of each magnesium double salt 
found at Stassfurt and indicate the ions found in a dilute aqueous 
solution of each. (Suggestion. — See under Potassium.) 

14. What weight of (a) crystallized magnesium sulphate and 
(6) crystallized zinc sulphate can be made from 275 gm. of the re- 
spective metals ? 

15. What weight of crystallized zinc chloride can be made from 
hydrochloric acid and 45 gm. of zinc oxide ? 

16. Calculate the simplest formulas from (a) Hg = 84.92, Cl = 
15.07; (b) Hg=: 73.8, CI -26.2. 

17. A solution of hydrochloric acid contains 39.1 per cent of HC1. 
What weight of zinc (95 per cent pure) is needed to liberate the 
hydrogen from 172 gm. of the solution? 

18. Calculate the solubility product of magnesium hydroxide 
(Mg(OH) 2 ) if the ionization is 80 per cent and the molar solubility 
is .00015. And of mercuric sulphide (HgS) if the ionization is 100 
per cent and the molar solubility is .02617. 



CHAPTER XXIX 
Aluminium 

Occurrence. — Aluminium (or aluminum), Al, does not 
occur free in nature, but its compounds are numerous and 
widely distributed. About 8 per cent of the earth's crust 
is aluminium. It is the most abundant metal and the third 
element in abundance in the earth's crust. Many common 
rocks are silicates of aluminium and other metals, e.g. feldspar 
and mica, which make up a large part of granite and gneiss. 
Clay and slate are mainly silicate of aluminium, which was 
formed by the decomposition of complex aluminium miner- 
als. Corundum and emery are more or less impure alumin- 
ium oxide (A1 2 3 ). Bauxite is an impure hydroxide of alu- 
minium (AI2O5H4). Cryolite is a fluoride of aluminium and 
sodium (Na 3 AlF 6 ). 

Metallurgy. — Aluminium is obtained from its purified 
oxide (A1 2 3 ) by electrolysis. An open iron vessel lined with 
carbon is made the cathode (Fig. 77). The anode consists 
of several carbon bars hung from a common copper rod, 
which can be lowered as the carbon is consumed. The bot- 
tom of the box is first covered with cryolite, the anode is 
lowered, and the box is then filled with cryolite. The cur- 
rent is turned on, and in its resisted passage through the 
cryolite enough heat is generated to melt the cryolite. Pure, 
dry aluminium oxide is now added, which is decomposed 
into aluminium and oxygen. The oxygen goes to the anode 
and unites with the carbon, forming carbon monoxide, which 
burns or escapes. The molten aluminium, which goes to the 

493 



494 



INORGANIC CHEMISTRY 



cathode, gradually collects on the bottom of the cell. The 
process is continuous, fresh aluminium oxide being added 



n + 




and the molten alumin- 
ium being drawn off at 
intervals. The cryolite is 
unchanged chemically. 

Properties. *— Alumin- 
ium is a lustrous, bluish 
white metal. It is very 
light compared with 
other common metals, 
since its specific gravity 
is only about 2.6; this 
value is one third that 
of iron. It is ductile and malleable, and is often sold 
in the form of wire and sheets; it must be annealed fre- 
quently during the hammering or drawing. It is a good 
conductor of heat and electricity. Its tensile strength is 
about as great as that of cast iron. It melts at about 658° 
C. and can be cast and welded, but not readily soldered so as 
to produce a permanent joint. Pure aluminium is only very 
slightly oxidized by air. Hydrochloric acid interacts readily 
with it, forming aluminium chloride and hydrogen, thus : — 



F*g. 77. — Apparatus for the manufacture of 
aluminium by the electrolysis of aluminium 
oxide. C, C, C is the iron box lined with 
carbon which serves as a cathode. A, A, 
etc., are carbon bars attached to the cop- 
per rod R. Connection is made with the 
cathode at D. 



2A1 


+ 


6HC1 


2 A1C1 3 


+ 


3H 2 


aminiu 


m 


Hydrochloric 
Acid 


Aluminium 
Chloride 




Hydrogen 



Sulphuric acid interacts feebly with aluminium, but nitric 
acid has no apparent effect. With a boiling solution of 
sodium or potassium hydroxide aluminium forms hydrogen 
and an aluminate, thus : — 



6 NaOH + 

Sodium Hydroxide 



2A1 = 2Na 3 A10 3 

Aluminium Sodium 

Aluminate 



+ 3H 2 

Hydrogen 



ALUMINIUM 495 

Uses. — The varied properties of aluminium adapt it to 
numerous uses. It is made into the metallic parts of military 
outfits, caps for jars, surgical instruments, cooking utensils, 
tubes, the fittings of boats, automobiles, and air ships, tele- 
phone receivers, scientific apparatus, parts of opera glasses 
and telescopes, the framework of cameras, stock patterns for 
foundry work, and hardware samples. Its attractive appear- 
ance has led to its extensive use as an ornamental metal, both 
in interior decorative work and in numerous small objects, 
such as trays, picture frames, hairpins, and combs. Alumin- 
ium leaf is used for decorating book covers and signs; a 
mixture of aluminium powder and. an adhesive oil is likewise 
used as a paint for steam pipes, lamp-posts, radiators, smoke- 
stacks, and other metal objects. Aluminium wire is used 
as a conductor of electricity. Large quantities of alumin- 
ium are used in the steel industry ; added to molten steel, 
the aluminium combines with gases and produces castings 
free from blow holes. Considerable is used in preparing cer- 
tain metals and in welding. Thus, if a mixture of chromium 
oxide and powdered aluminium is ignited at one point by 
a special device, the reduction thus initiated proceeds rapidly 
throughout the mixture and the intense heat thereby gener- 
ated fuses the chromium, which can be removed from the 
crucible subsequently as a coherent mass ; the aluminium 
oxide likewise melts and separates from the metal. The 
equation for the chemical change is — 



Cr 2 3 + 2 Al = 


2 Cr + A1 2 3 


Chromium Aluminium 


Chromium Aluminium 


Oxide 


Oxide 



Other metals hitherto rare or expensive can be similarly pre- 
pared. If a mixture of ferric oxide (Fe 2 3 ) and powdered 
aluminium is ignited, molten steel at a temperature of about 
3000° C. is produced. By means of a special apparatus the 
molten steel can be conducted to a joint or crack ; this pro- 



496 INORGANIC CHEMISTRY 

cess is used in welding iron rails and repairing fractures. 
These mixtures of aluminium and oxides are called " thermit," 
and* the method is known as the Goldschmidt or alumino- 
thermic method. 

Alloys. — The alloy of aluminium and copper — alumin- 
ium bronze — has been described. (See Alloys of Copper.) 
Magnalium contains from 75 to 90 per cent of aluminium, the 
rest being magnesium ; it is used in scientific instruments, 
e.g. as balance beams. 

Aluminium Oxide, or alumina, A1 2 3 , is the only oxide of 
aluminium. Pure crystalline varieties of native alumina are 
known as corundum, while the impure forms are called 
emery ; they are very hard substances, corundum being next 
below the diamond in the scale of hardness. Emery was 
formerly used as an abrasive, but it has been largely displaced 
by carborundum (see page 385). The transparent colored 
kinds of corundum are prized as gems (see below). 

Alumina is prepared as a white powder by heating the 
metal in the air or by heating the hydroxide. The product 
obtained by heating the hydroxide to a low temperature 
interacts with acids but the intensely heated oxide resembles 
the crystallized varieties in being almost insoluble in acids. 
The native oxide is converted into an aluminate and a sul- 
phate respectively by fusing with a caustic alkali (e.g. potas- 
sium hydroxide) and with acid potassium sulphate (HKS0 4 ). 
Alumina melts in the oxyhydrogen flame, in the electric fur- 
nace, and during the reduction of oxides by aluminium (see 
Goldschmidt method above). 

Compact crystalline alumina is manufactured from bauxite 
by an electrothermic process. It resembles native corundum 
and is used as an abrasive. This artificial alumina is known 
in trade as "alundum." 



ALUMINIUM 497 

Gems containing Aluminium. — Corundum (A1 2 3 ) has long been 
found in crystals in Ceylon, Siam, Burma, and other places in the 
Orient. The color is due to traces of impurities, usually oxides of 
metals. The sapphire is blue and the ruby is red. The Oriental 
topaz is yellow, the Oriental amethyst is purple, and the Oriental 
emerald is green. Sapphires, rubies, and similar gems are now made 
by melting aluminium oxide (with or without coloring matter) in an 
oxyhydrogen flame. Spinels are complex compounds of alumin- 
ium. The typical or ruby spinel is magnesium aluminate (MgAl 2 04). 
It resembles the true ruby in color. Other spinels differ from 
the ruby spinel both in color and in composition. Turquoise is 
a complex aluminium phosphate containing traces of copper. It 
has a beautiful robin's-egg-blue color, is compact, and may be worked 
into various shapes. Formerly turquoise came almost exclusively 
from Persia, but now New Mexico meets most demands. Topaz is 
a complex aluminium silicate containing fluorine, usually has a pale 
yellow color, and is found in many localities. Emerald is the most 
precious gem next to diamond and ruby. It is an aluminium silicate 
containing the rare element beryllium. The finest specimens have a 
deep emerald green color and come from Colombia, South America. 
Garnet is a complex silicate of aluminium and another metal, espe- 
cially calcium, magnesium, iron, or manganese. The kind used as a 
gem has a deep red color and is rather abundant. 

Aluminium Hydroxide, Al(OH) 3 , is a white, jellylike solid 
formed by the interaction of an hydroxide and the solution 
of an aluminium salt, e.g. : — 

AICI3 + 3NH 4 OH = Al(OH) 8 + 3NH 4 C1 

Aluminium Ammonium Aluminium Ammonium 

Chloride Hydroxide Hydroxide Chloride 

It has weak acid and basic properties, the latter, however, 
being the stronger. Its acid property is shown by the forma- 
tion of soluble saltlike compounds called aluminates upon 
the addition of an excess of sodium or potassium hydroxide, 
thus : — 



Al(OH) s 


+ 


3 NaOH = 


= Na 3 A10 3 


+ 


3H 2 


Aluminium 




Sodium 


Sodium 




Water 


Hydroxide 




Hydroxide 


Aluminate 







498 INORGANIC CHEMISTRY 

The feeble basic property of aluminium hydroxide is illus- 
trated by the fact that it does not form salts with such weak 
acids as carbonic, hydrosulphuric (H 2 S), and sulphurous 
(H 2 S0 3 ). Thus, when sodium carbonate or ammonium 
sulphide is added to the solution of an aluminium salt, alu- 
minium hydroxide (not the carbonate or sulphide) is pre- 
cipitated, e.g. : — 

2 AICI3 + 3 (NH 4 ) 2 S + 6 H 2 = 2 Al(OH) 3 + 6 NH 4 C1 + 3 H 2 S 



Alumin- 


Ammo- 


Water 


Alumin- 


Ammo- 


Hydro- 


ium 


nium 




ium 


nium 


gen 


Chloride 


Sulphide 




Hydroxide 


Chloride 


Sulphide 



Like all bases, however, aluminium hydroxide interacts with 
strong acids and thereby forms salts, thus : — 



Al(OH) 3 


+ 3 HCl 


= A1C1 8 


+ 


3H 2 


Aluminium 


Hydrochloric 


Aluminium 




Water 


Hydroxide 


Acid 


Chloride 







Commercial aluminium hydroxide is prepared by roasting 
bauxite or cryolite with sodium carbonate, extracting the 
resulting aluminate with water, and precipitating the hy- 
droxide by passing carbon dioxide into the solution; the 
last operation may be represented by the equation : — 

2Na 3 A10 3 + 3C0 2 + 3 H 2 = 2A1(0H) 3 + 3 Na 2 C0 3 

Sodium Carbon Water Aluminium Sodium 

Aluminate Dioxide Hydroxide Carbonate 

The aluminium hydroxide is dried and sold in the form of a 
white powder. There are several native aluminium hydrox- 
ides. Bauxite (A1 2 5 H 4 or A1 2 3 . 2 H 2 0) contains ferric oxide 
(Fe 2 3 ) as an impurity. It resembles clay in texture and 
color. The vast deposits found at Baux, in southern France, 
furnish much of the raw material for the manufacture of 
aluminium, though some is obtained from our Southern 
states. Hydrargyllite (Al(OH) 3 ) and diaspore (A10 2 H) are 
found in relatively small quantities. 



ALUMINIUM 499 

Aluminates have been described in the foregoing para- 
graphs. They are compounds in which aluminium acts as 
an acid element, corresponding in this respect to the zincates. 
Aluminates are soluble to some extent in water and such 
solutions have an alkaline reaction owing to hydrolysis; the 
equation for the hydrolysis is — 

Na 3 A10 3 + 3 H 2 = Al(OH) 8 + 3 NaOH 

the alkaline reaction being caused by the hydroxyl ions lib- 
erated by the ionization of the sodium hydroxide. No 
aluminate is formed by adding an excess of ammonium hy- 
droxide to aluminium hydroxide; this property is some- 
times used to distinguish aluminium from zinc, since zinc 
hydroxide (a similar compound) forms a soluble complex 
compound by the interaction with ammonium hydroxide. 

Aluminium Sulphate, A1 2 (S0 4 ) 3 . 18 H 2 0, is a white, crystal- 
line solid. The commercial salt has a variable composition; 
if pure, it dissolves readily and completely in water. It is 
extensively used as a mordant in dyeing (see below), as an 
ingredient of the size put upon paper to prevent the ink from 
spreading, and in purifying water. 

A solution of aluminium sulphate has an acid reaction on 
account of hydrolysis; the equation for the hydrolysis is — 

A1 2 (S0 4 ) 3 + 6 H 2 = 2 Al(OH) 3 + 3 H 2 S0 4 

the acid reaction being due to the hydrogen ions liberated 
by the ionization of the sulphuric acid. 

Aluminium sulphate is prepared from pure clay, bauxite, or 
cryolite. If clay or bauxite is heated with sulphuric acid and then 
allowed to cool, the product is impure aluminium sulphate, known as 
"alum cake," or if much iron is present, as u alumino-ferrie cake"; it 
is used to purify sewage and for other purposes where iron and the 
other impurities are harmless. Purer aluminium sulphate is prepared 
by heating bauxite with sodium carbonate, extracting the sodium 
aluminate with water, and precipitating the aluminium as the hydrox- 



500 INORGANIC CHEMISTRY 

ide with carbon dioxide gas; the relatively pure hydroxide is then 
changed into sulphate by treatment with sulphuric acid. The 
product, known as " concentrated alum/' has the composition ex- 
pressed by the formula A1 2 (S04)3 . 20 H 2 0, though separate crystals 
may contain only eighteen molecules of water of crystallization. By 
boiling cryolite with milk of lime, the sodium aluminate thereby 
formed may be changed into " concentrated alum," as described above. 
About 50,000 tons of "concentrated alum" are annually produced 
in the United States. 

Alum. — When solutions of aluminium sulphate and po- 
tassium sulphate are mixed and concentrated by evapora- 
tion, transparent, colorless, glassy crystals are deposited. 
This solid is potassium alum or simply alum. It has the 
composition represented by the formula, K 2 A1 2 (S0 4 ) 4 . 24 H 2 0, 
or K 2 S0 4 , A1 2 (S0 4 ) 3 • 24 H 2 0, and is sometimes called a double 
salt. It is the type of a class of similar salts called alums, 
which can be formed by mixing the solution of a sulphate of 
a trivalent metal {e.g. Al, Cr, Fe) with the solution of a sul- 
phate of a univalent metal {e.g. K, Na, NH 4 ). Alums are 
rather soluble in water, and their solutions have an acid 
reaction owing to hydrolysis. (See Aluminium Sulphate.) 
They crystallize as octahedrons and contain twenty-four 
molecules of water of crystallization. When heated, alums 
lose their water of crystallization and some sulphur trioxide 
and fall to a white powder or porous mass known as burnt 
alum. Potassium alum is the most common, but ammonium 
and sodium alums are manufactured and used. Sodium 
alum is an ingredient of some baking powders. Burnt alum 
finds application as a medicine. Alum has been largely dis- 
placed by " concentrated alum/ 7 since aluminium sulphate 
has the same general properties; but the real alum is still 
used to some extent in dyeing and printing cloth, in tanning 
and paper making, in purifying water and sewage, as a medi- 
cine, for hardening plaster, in making wood and cloth fire- 
proof, and in preparing other aluminium compounds. 



ALUMINIUM 501 

Alum was known to the ancients, who used it in dyeing and tan- 
ning, and as a medicine. It was first manufactured in Europe, about 
the thirteenth century, from native alunite, which is an impure sul- 
phate of aluminium, potassium, and iron. Alunite and alum slates 
or shales are now used to some extent, but most of the alum is made 
from bauxite. 

Not all alums contain aluminium. As stated in the preceding 
paragraph, this metal may be replaced by iron, chromium, manganese, 
or any metal having the valence of three. Hence the general formula 
of an alum is M 2 (SO^s • X2SO4 . 24 H 2 0, in which M may be aluminium, 
iron, chromium, etc., and X a metal (or group) like potassium, sodium, 
ammonium. 

Alums are double salts, i.e. crystalline compounds of two or more 
normal salts. Their dilute aqueous solutions contain the ions of the 
separate salts and have properties which indicate the double nature, 
so to speak, of the salt in question. (Compare the Stassfurt salts.) 

Alums and other aluminium salts are used as mordants in 
dyeing and calico printing. Some dyes must be fixed in the 
fabric by a metallic substance, otherwise the color could be 
easily removed. The cloth to be dyed or printed is impreg- 
nated or printed with a mordant, and then heated or treated 
with some substance to change the mordant into an insoluble 
compound. The mordanted cloth is next passed through a 
vat containing the solution of the dye, which unites chemi- 
cally or mechanically (perhaps both) with the metallic 
compound, forming a colored compound. The latter is 
called a "lake"; it is relatively insoluble, and cannot be 
easily washed from the cloth, i.e. it is a fast color. Alumin- 
ium acetate or "red liquor/ 7 aluminium sulphate, and 
sodium aluminate, besides alum, are used as mordants for 
cotton, linen, and w r ool. The use of aluminium salts as 
mordants depends upon the fact that they hydrolyze readily; 
it is the resulting aluminium hydroxide, therefore, which is 
the effective metallic compound in dyeing. 

Cryolite is a white, glassy, crystalline solid. It often 
resembles clouded ice, and its name means "ice stone. " 



502 INORGANIC CHEMISTRY 

Its composition corresponds to the formula Na 3 AlF 6 (or 
A1F 3 . 3 NaF). Small fragments melt easily, even in a candle 
flame, and color the Bunsen flame yellow. The only local- 
ity where it is found in commercial quantities is southern 
Greenland, which yields annually about 10,000 tons. It is 
used not only in manufacturing aluminium, but as a source 
of alum and aluminium hydroxide, pure sodium carbonate 
and hydroxide, hydrofluoric acid, and fluorides. 

Aluminium Chloride when pure is a white powder, but it is 
often a yellowish, crystalline mass (A1C1 3 . 6 H 2 0). It is pre- 
pared by heating powdered aluminium in chlorine, or by 
passing chlorine over a heated mixture of aluminium oxide 
and carbon. Exposed to the air, it absorbs moisture and 
gives off fumes of hydrochloric acid. It dissolves in water 
with evolution of heat, and if the solution is heated, hydro- 
chloric acid is expelled, owing to the hydrolysis of the chloride, 
thus : — 

AICI3 + 3 H 2 = 3 HC1 + Al(OH) 3 

Aluminium Water Hydrochloric Aluminium 

Chloride Acid Hydroxide 

This salt is used in organic chemistry. 

Tests for Aluminium. —7 When a compound of aluminium 
is heated on charcoal with a blowpipe, cooled, moistened 
with cobaltous nitrate solution, and then reheated, the mass 
finally turns a beautiful blue color. Sodium hydroxide pre- 
cipitates white gelatinous aluminium hydroxide, which dis- 
solves in an excess of the alkali (NaOH), owing to the forma- 
tion of soluble sodium aluminate. Aluminium hydroxide 
is insoluble in an excess of ammonium hydroxide (distinc- 
tion from zinc hydroxide). 

Miscellaneous. — It is evident from the preceding para- 
graphs that aluminium, like zinc, acts both as a metal and a 



ALUMINIUM 503 

non-metal. Many soluble aluminium compounds yield color- 
less aluminium ions (Al + + + ). 

The atomic weight of this element is 27.1 and the valence 
is three. 

Clay is a more or less impure aluminium silicate, formed by the 
slow decomposition of rocks containing aluminium, especially feld- 
spar. Pure feldspar is- a silicate of aluminium and sodium or potas- 
sium. The products of its decomposition are chiefly an insoluble 
aluminium silicate and a soluble alkaline silicate. The latter is 
washed away. The aluminium silicate which remains is pure clay 
or kaolin. The latter is really a hydrous silicate, having the com- 
position corresponding to the formula Al 2 Si207 . 2 H 2 0. The com- 
position of clay varies, because it is seldom formed from pure feldspar. 
Pure kaolin is a white powdery substance, but most kaolin contains 
particles of mica and quartz. Ordinary clay contains many impuri- 
ties, e.g. carbonates of calcium and magnesium, quartz, and iron 
compounds. All kinds of clay become plastic when wet and can be 
molded into various objects which shrink on drying but retain their 
general form; if heated, the dried clay does not melt (except at a 
very high temperature) but becomes a permanently hard mass. 
These properties (plasticity when wet and hardness when heated) 
make clay a most serviceable substance. For ages it has been made 
into useful and ornamental articles which may be roughly put into 
three comprehensive classes — porcelain, pottery, and materials of 
construction. The classes as well as the varieties in each class are 
the result of differences in quality and proportions of raw material 
and in method of heating. The subvarieties merge into each other. 

Porcelain (china or chinaware) is the finest clay product. It is 
made by fusing a mixture of very pure kaolin, fine sand, and a more 
fusible substance, usually feldspar, though sometimes chalk and 
gypsum are also used. The fused mass when cool is hard, dense, 
white, and translucent (if thin); it is often called "biscuit" or "bis- 
cuit ware." Although not very porous, its surface is glazed, partly 
for protection, partly for ornament. This is done by coating the 
ware with a thin mixture similar to that used for making the porcelain 
but more easily fused, and then heating again so that the glaze will 
penetrate the clay. Pottery is a very large class and includes colored, 
white, glazed, and unglazed ware. The raw material is not as pure 



504 INORGANIC CHEMISTRY 

as that used for porcelain, nor is the mixture heated to such a high 
temperature; the product, therefore, is rather coarse, opaque, heavy, 
and porous. The finer varieties, such as crockery, art pottery, and 
some kinds of stone ware, are glazed like porcelain. The coarser 
varieties, such as jugs and domestic utensils, are glazed by throwing 
salt into the oven just before the firing (i.e. the baking or heating) is 
over; the sodium in the sodium chloride forms a fusible aluminium 
silicate, which coats the surface. Unglazed pottery is familiar under 
different names, e.g. flower pots, tiles, terra cotta, and clay tobacco 
pipes. Materials of construction made from clay include some grades 
of stoneware, bricks, pipes, drain tile, etc. The raw material is usu- 
ally impure and the firing is done at a low temperature. The product 
varies with the quality of the clay. Thus, ordinary clay containing 
iron compounds gives coarse red brick, while clay containing consider- 
able silica gives bricks which withstand high temperature and are 
called fire-clay bricks. Pipe and drain tile, both glazed and porous, 
are used to convey water, sewage, and other fluids, and as a conduit 
for underground electric wires, especially wire cables. 

Problems and Exercises 

1. What is the per cent of aluminium in (a) cryolite (NasAlF 6 ), 
(6) turquoise (A1 2 P 2 8 . H 6 A1 2 6 . 2 H 2 0), (c) corundum (A1 2 8 ), (d) 
aluminium hydroxide ? 

2. What volume of oxygen at 15° C. and 760 mm. is needed to 
change 5 gm. of aluminium to aluminium oxide (Al 2 Os) ? 

3. If 6.917 gm. of aluminium bromide (AlBrs) require 8.4429 
gm. of silver to precipitate all the bromine, what is the atomic weight 
of aluminium? (Assume Ag = 107.88 and Br = 79.92.) 

4. Write the formulas of the following compounds by applying 
the principle of valence and calculate the per cent of aluminium in 
each: Aluminium sulphide, aluminium phosphate, aluminium acetate, 
potassium aluminate. 

5. Compare the corresponding compounds of aluminium and 
zinc. 

6. Discuss (a) " aluminium acts as a metal and a non-metal," 
and (6) hydrolysis of aluminium compounds. 



CHAPTER XXX 
Tin and Lead — Cerium and Thorium 

Tin and lead are familiar metals. They have similar and 
useful properties, which give these metals and their com- 
pounds numerous applications. 

Tin 

Occurrence of Tin. — Metallic tin is rarely if ever found. 
Tin dioxide (cassiterite or tin stone, Sn0 2 ) is the only avail- 
able ore. It is not widely distributed, but large deposits 
are found in England (at Cornwall), Germany (in Bohemia 
and Saxony), Australia, Tasmania, and the East Indian 
Islands, especially Banca and Billiton. A small quantity is 
found, but not mined, in the United States. 

Tin is one of the oldest known metals. It is mentioned in the 
Pentateuch and was probably obtained long before the Christian 
era by the Phoenicians from the British Isles, which were called 
Cassiterides (from the Greek word kassiteros, meaning tin). Many 
ancient bronzes contain tin. The alchemists called it Jupiter, and 
used the metal and its compounds. 

The Latin word stannum gives the symbol Sn and the terms 
stannous and stannic. 

Metallurgy of Tin. — If the tin ore contains sulphur or 
arsenic, these impurities must be removed by roasting. 
The tin oxide is then reduced by heating it with coal in a 
reverberatory furnace; the simplest equation for this change 
is — 

Sn0 2 + C = Sn + C0 2 

Tin Carbon Tin Carbon 

Dioxide Dioxide 

505 



506 INORGANIC CHEMISTRY 

The molten tin, which collects at the bottom of the furnace, 
is drawn off and cast into bars or masses, which are often 
called block tin. Usually it is purified by melting it slowly 
on a hearth, inclined so that the more easily melted tin will 
flow down the hearth and leave the metallic impurities be- 
hind. This tin may be further purified by stirring the molten 
metal with a wooden pole or by holding billets of wood 
beneath its surface. The impurities, which are oxidized by 
the escaping gases, collect as a scum on the surface and are 
removed. 

A small amount of tin is obtained by treating rejected 
tin plate or scrap tin with chlorine or some other active 
dissolving chemical. 

Properties of Tin. — Tin is a white, lustrous metal. It 
is soft and malleable, and can be readily cut and hammered. 
It is softer than zinc but harder than lead. Its specific 
gravity is 7.3. Tin may be obtained in the crystalline 
form, and when a piece of such tin is bent it makes a crac- 
kling sound, which is caused by the friction of these crystals 
upon one another. It melts at about 232° C, and # when 
heated to a higher temperature it burns, forming white 
tin oxide (Sn0 2 ). The physical properties of tin, like those 
of zinc, vary with the temperature. One property of 
tin is rather striking. If kept at a low temperature for 
some time, ordinary tin changes slowly into a gray powder 
having a specific gravity of about 5.8. Experiment shows 
that this transformation begins at 18° C, and this tempera- 
ture is called the transition point. That is, ordinary tin is 
stable only above 18° C. ; below 18° C. it is unstable and may 
form gray tin. Sometimes this " tin disease," as it might be 
called, attacks the pipes of church organs. The appearance 
of a sheet of tin affected by " tin disease " is shown in Fig- 
ure 78. Concentrated hydrochloric acid changes tin into 



TIN 



507 



stannous chloride (SnCl 2 ), hot concentrated sulphuric acid 
converts it into stannous sulphate (SnS0 4 ), while commer- 
cial nitric acid transforms it into a white solid (metastannic 
acid (H 2 Sn0 3 ) 5 ). With sodium hydroxide tin forms sodium 




^^^^^^^^P^ 



Fig. 78. — Sheet of tin affected by "tin disease " (enlarged one and one 

half times). 

metastannate (Na 2 Sn0 3 ). Zinc precipitates tin from its so- 
lutions as a grayish black, spongy mass, which is sometimes 
filled with bright scales. This is due to the fact that zinc 
precedes tin (by several places) in the electromotive series 
of the metals. (See Chapter XXVI, Electromotive Series of 
Metals.) 



Uses of Tin. — Tin is so permanent in air, water, weak 
acids (like vinegar and fruit acids), and alkalies that it is 
extensively used in making scientific apparatus and as a 
protective coating for metals. Condensing pipes in stills 
are often made of tin. Ordinary tinware is sheet iron 
coated with tin. The tin plate (sheet tin, or simply "tin") 
is made by dipping very clean sheet iron into molten tin. 
Tacks, nails, and many small iron objects are similarly 
tinned. Copper coated with tin is made into vessels for 
cooking, and brass coated with tin is made into pins. Large 



508 INORGANIC CHEMISTRY 

quantities of tin plate are used to cover roofs. Tinned iron 
does not rust until the tin is worn off and the iron exposed, 
and then the rusting proceeds rapidly. Tin is also ham- 
mered into thin sheets called tin foil, though much tin foil 
contains lead. Many useful alloys contain tin as an essen- 
tial ingredient. 

Alloys of Tin containing a minor percentage of tin are 
bronze, gun metal, speculum metal, type metal, anti-fric- 
tion metals, and fusible alloys. Britannia metal contains 
about 90 per cent tin, 8 per cent antimony, and the rest 
mainly copper. It is a white metal, and was formerly made 
into tableware. White metal contains less tin and more 
antimony than Britannia, though the composition varies. 
It resembles Britannia. The harder varieties of white metal 
are used as parts of machinery, and the softer kinds are made 
into ornaments and cheap jewelry. Pewter and solder 
contain varying proportions of tin and lead. Plumbers' 
solder, or soft solder, is about one third tin and two thirds 
lead; it is harder than either constituent, but melts at a 
lower temperature. Tin amalgam is sometimes used to 
coat mirrors. 

Compounds of Tin. — Tin forms two series of compounds, 
the stannous and the stannic. The valence of tin is two in 
the stannous compounds and four in the stannic. Stannic 
oxide (Sn0 2 ) has already been mentioned as the chief ore 
of tin and the product formed when tin is burned. The 
artificial oxide is faint yellow when hot and white when cold. 
The native oxide is a brown or black, lustrous, (often) crys- 
talline solid. Irregular pebbles called stream tin occur 
in some localities near rivers. Stannous chloride (SnCl 2 ) 
is formed by the interaction of hydrochloric acid and tin. 
From the concentrated solution a greenish salt crystallizes 
(SnCl 2 . 2H 2 0), known as the tin crystals or salt of tin. 



TIN 509 

Stannous chloride passes readily into stannic chloride 
(SnCl 4 ) when added to mercuric chloride solution. The 
simplest equation for this change is — 

SnCl 2 + 2HgCl 2 = SnCl 4 + 2 HgCl 

Stannous Mercuric Stannic Mercurous 

Chloride Chloride Chloride Chloride 

By an extension of the simplest idea of oxidation and reduc- 
tion, the stannous chloride in this change is said to be 
oxidized to stannic chloride and the mercuric chloride to be 
reduced to mercurous chloride. An excess of stannous 
chloride reduces the white mercurous chloride to a gray 
precipitate of finely divided mercury; this reaction is used 
as a test for tin. (Compare tests for Mercury, Chapter 
XXVIII.) Stannous chloride is often used as a reducing 
agent and as a mordant in calico dyeing and printing. Stan- 
nic chloride (SnCl 4 ) is a colorless, fuming liquid; it forms 
a crystalline hydrate (SnCl 4 . 5H 2 0), known commercially 
as oxymuriate of tin, which is used as a mordant. Am- 
monium stannic chloride ((NH 4 ) 2 SnCl 6 ), or "pink salt," 
is also used as a mordant. Tin mordants produce brilliant 
colors. Sodium stannate (Na 2 Sn0 3 . 3 H 2 0) is extensively 
used to prepare cotton cloth for printing. With hydrogen 
sulphide, stannous compounds form brown stannous sulphide 
(SnS),and stannic compounds form yellow stannic sulphide 
(SnS 2 ) ; both sulphides dissolve in ammonium polysulphide, 
owing to the formation of soluble sulpho-salts of tin. 

Miscellaneous. — Tin, like zinc and aluminium, acts both 
as a metal and a non-metal. Thus, there are the stannous 
and stannic salts and the stannates. Solutions of stannous 
salts contain stannous ions (Sn ++ ). 

The atomic weight of tin is 119.0, and the valence, as 
stated above, is two in stannous and four in stannic com- 
pounds, 



510 INORGANIC CHEMISTRY 



Lead 

Occurrence of Lead. — Metallic lead is occasionally found 
in small quantities. The most abundant ore is lead sul- 
phide (galena, PbS). Other native compounds are the car- 
bonate (cerussite, PbC0 3 ), the sulphate (anglesite, PbS0 4 ), 
and the phosphate (pyromorphite, PbsOKPO^). 

Lead has been used by civilized people since the dawn of history. 
The Chinese have used it for ages to line chests in which tea is stored 
and transported. The Romans called it plumbum nigrum, i.e. black 
lead. The symbol Pb comes from plumbum. 

Lead ores are found in the United States mainly in the Middle 
West, Colorado, Idaho, and Utah. 

Metallurgy of Lead. — Lead is obtained from galena by 
several processes. (1) Ores rich in lead are roasted in a 
reverberatory furnace until a part of the sulphide is changed 
into lead oxide and lead sulphate. Thus : — 

2 PbS + 3 2 = 2 PbO + 2 S0 2 

Lead Oxygen Lead Sulphur 

Sulphide Oxide Dioxide 

PbS + 2 2 = PbS0 4 

Lead Oxygen Lead 

Sulphide Sulphate 

The air is then excluded and the temperature raised ; the 
mixture interacts thus : — 

2 PbS + PbS0 4 + 2 PbO = 5 Pb + 3 S0 2 

Lead Lead Lead Lead Sulphur 

Sulphide Sulphate Oxide Dioxide 

(2) Ores poor in lead are sometimes reduced by roasting 
with iron, the equation for the reaction being : — 

PbS + Fe = Pb + FeS 

Lead Iron Lead Iron 

Sulphide Sulphide 

(3) Lead ores rich in silver are roasted and then reduced with 
coal (or coke), limestone, and iron ore. 



LEAD 511 

Lead is refined by heating it to oxidize most of the copper, 
arsenic, and antimony, and then treating the alloy by the 
Parkes process (page 453). In an electrolytic process the 
cathode is a sheet of pure lead, the anode is a plate of impure 
lead, and the electrolytic solution is a mixture of lead fluo- 
silicate (PbSiF 6 ) and gelatin; pure lead is deposited on the 
cathode and most of the other metals remain attached to 
the remna t of the anode, from which they are subsequently 
recovered, especially the gold and silver. 

Properties of Lead. — Lead is a blue-gra}^ metal. When 
scraped or cut, it has a brilliant luster, which soon dis- 
appears, owing to the formation of a film of oxide or other 
lead compound. This coating protects the lead from further 
change. It is a soft metal, and can be scratched with the 
finger nail. It discolors the hands, and when drawn across 
a rough surface it leaves a black mark. For this reason it 
is sometimes erroneously called black lead. (See Graphite.) 
Lead is not tough nor very ductile, though it can be made into 
wire, pressed (while soft) into pipe, and rolled into sheets. 
It is a heavy metal, its specific gravity being about 11.4; 
with the exception of mercury, it is the heaviest of the familiar 
metals. It melts at 327° C, or about 100° higher than tin 
and 100° lower than zinc. Lead, when heated strongly 
in the air, changes into lead monoxide (litharge, PbO) ; at 
higher temperatures (about 350° C. and above) the tetroxide 
(Pb 3 4 ) is formed. Hydrochloric and sulphuric acids exert 
only a slight chemical action upon compact lead. (See Lead 
Sulphate.) Nitric acid changes it into lead nitrate (Pb(N0 3 )2). 
Acetic acid (or vinegar) and acids from fruits and vegetables 
change it into soluble, poisonous compounds ; hence cheap 
tin-plated vessels, which sometimes contain lead, should 
never be used in cooking. Zinc and iron precipitate lead from 
its solutions as a grayish mass, which often has a beautiful 
treelike appearance. 



512 INORGANIC CHEMISTRY 

This displacement, as in the case of tin, is due to the fact 
that lead is lower than zinc and iron in the electromotive 
series. 

Lead in Drinking Water. — Lead is slowly changed into 
soluble compounds by water containing free oxygen, carbon 
dioxide, ammonia, nitrates, or chlorides. But water con- 
taining sulphates or carbonates forms an insoluble coating 
on the lead, thus protecting it from further action. All 
lead salts are poisonous; and if taken into the system they 
slowly accumulate and ultimately cause serious and danger- 
ous illness. Water suspected of attacking lead should 
never be drunk after it has been standing very long in lead 
pipes, but should be allowed to flow until the pipe has been 
filled with fresh water. It is sometimes safer to substitute 
an iron or block tin pipe for the customary lead service pipe. 

Uses of Lead. — Lead on account of its plasticity when 
warm is extensively made into pipe, which can be easily 
cut, bent, and united (by solder). Lead pipe is not only 
used to convey water to and from parts of buildings, but 
as a sheath for electric wires, both overhead and under- 
ground. In the form of sheets it is used to cover roofs 
and to line sinks, cisterns, and the cells employed in some 
electrolytic processes. The lead chambers and evaporating 
pans used in manufacturing sulphuric acid are made of 
sheet lead. Shot and bullets are lead (alloyed with a little 
arsenic) . 

The Alloys of Lead are important. Type metal contains 
70 to 80 per cent lead; the other constituents are tin and 
antimony. The alloy is harder than the lead itself and 
expands on cooling, thereby making the face of the type 
sharp and hard. Solder, pewter, and fusible alloys contain 
lead as an essential constituent. (See these Alloys.) 



LEAD 513 

Lead Oxides. — There are three important oxides. Lead 
monoxide (PbO) is a yellowish powder known as massicot, 
or a buff-colored crystalline mass called litharge. It is 
formed by heating lead above its melting point in a current 
of air. It is made this way, though considerable is obtained 
as a by-product in separating silver from lead. (See Cupella- 
tion.) Large quantities are used in preparing certain oils 
and varnishes, flint glass, other lead compounds, and as a 
glaze for pottery. Lead tetroxide (red lead or minium, 
Pb 3 4 ) is a red powder, varying somewhat in color and 
composition. It is prepared by heating lead (or lead mo- 
noxide) to about 350° C. It is used in making flint glass. 
Pure grades are made into artists 7 paint, but the cheap 
variety is used to paint structural iron work (bridges, gas- 
ometers, etc.), hulls of vessels, and agricultural implements. 
A mixture of red lead and oil is used in plumbing and gas 
fitting to make joints tight. Orange mineral has the same 
composition as red lead, although its color is lighter; its 
uses are the same. Lead dioxide (lead peroxide, Pb0 2 ) is 
a brown powder formed by treating lead tetroxide with 
nitric acid or by the action of chlorine on an alkaline solution 
of lead acetate. It is used in storage batteries. 

Lead Carbonate, PbC0 3 , is found native as the trans- 
parent, crystalline mineral cerussite. It is obtained as a 
white powder by adding ammonium carbonate solution to 
lead nitrate solution. Sodium or potassium carbonate, 
however, forms basic lead carbonates, whose composition 
depends upon the temperature. The most important of 
these basic carbonates usually has the composition corre- 
sponding to the formula Pb3(OH) 2 (C0 3 )2, and is known as 
white lead. It is a heavy white powder which mixes well 
with linseed oil, and is used extensively as a white paint 
and as the basis of many colored paints. Its value as a 



514 INORGANIC CHEMISTRY 

paint is largely due to its superior covering power, i.e. a 
very thin layer produces a perfectly white, opaque surface. 
In recent years other substances have been mixed with or 
substituted for white lead, e.g. zinc oxide, barium sulphate, 
and lithophone. These solids are white, do not darken in 
air (as white lead does), and often improve the paint in other 
ways. 

White lead is manufactured by several processes. The Dutch 
process is the oldest, having been used as early as 1622. It is essen- 
tially the same to-day, though many details have been improved. 
Perforated disks of lead are put in earthenware pots which have a 
separate compartment at the bottom containing a weak solution of 
acetic acid (about as strong as vinegar). These pots are arranged 
in tiers in a large building, and spent tan bark is placed between each 
tier. The building is now closed, except openings for the entrance 
and exit of air and steam. The heat volatilizes the acetic acid, 
which changes the lead into a lead acetate. The tan bark ferments 
and liberates carbon dioxide, which changes the lead acetate into 
basic lead carbonate. The whole operation requires from sixty to one 
hundred days. Other processes, much the same as the Dutch process, 
are used, their chief aim being to lessen the time of manufacture. 

Lead Sulphide, PbS. — Native lead sulphide is the mineral 
galena, the chief ore of lead. It resembles lead in appear- 
ance, but is harder and is usually crystallized as cubes, 
octahedrons, or their combinations. It has a perfect cubic 
cleavage, i.e. it breaks into cubes or fragments more or less 
rectangular. It is easily changed into lead by heating it 
alone or with sodium carbonate on charcoal. Lead sulphide 
as prepared in the laboratory is a black solid. Black lead 
sulphide is readily precipitated from a lead salt solution 
by hydrogen sulphide. Its formation is a test for lead. It 
is changed into lead chloride by concentrated hydrochloric 
acid and into lead sulphate by concentrated nitric acid. 
Additional tests for lead are the formation of the sulphate 
and chromate, as described in the next paragraph. 



CERIUM AND THORIUM 515 

Other Compounds of Lead, which are important, are the 
chloride, sulphate, nitrate, chromate, and acetate. Lead 
chloride (PbCl 2 ) is a white solid formed by adding hydro- 
chloric acid or a soluble chloride to a cold solution of a lead 
salt. It dissolves in hot water. Lead sulphate (PbS0 4 ) is a 
white solid formed by adding sulphuric acid or a soluble 
sulphate to a solution of a lead salt. It is very slightly 
soluble in water, but soluble in sulphuric acid, hence crude 
sulphuric acid often contains lead sulphate. Lead nitrate 
(Pb(N0 3 ) 2 ) is a white crystalline solid formed by dissolving 
lead (or better, lead monoxide) in nitric acid. When heated, 
it decomposes into lead oxide (PbO), nitrogen peroxide, 
and oxygen. Lead chromate (PbCr0 4 ) is a yellow solid 
formed by adding a solution of a lead compound to a 
solution of potassium chromate or potassium dichromate. 
It is sometimes called " chrome yellow." Lead acetate 
(Pb(C 2 H 3 02)2) is a white crystalline solid formed by the 
action of acetic acid upon lead or lead oxide (PbO). It 
is very soluble in water and is often called " sugar of lead." 

Miscellaneous. — Aqueous solutions of lead compounds 
may contain several kinds of ions. The common com- 
pounds, such as lead chloride, lead nitrate, and lead acetate, 
yield a colorless ion (Pb ++ ). Other compounds yield ions 
containing oxygen as well as complex ions. Compounds in 
which lead acts as a non-metal are known, e.g. sodium 
plumbate (Na 2 Pb0 3 ). 

The atomic weight of lead is 207.10. The valence is two 
in many of its compounds; it is four in lead dioxide (Pb0 2 ) 
and plumbates. 

Cerium and Thorium 

Cerium (Ce) and Thorium (Th) are members of a family 
of rare elements in the same periodic group as tin and lead. 
They are constituents of rare minerals, but their compounds 



516 INORGANIC CHEMISTRY 

are prepared from a complex mineral substance named 
monazite sand. The oxides of thorium and cerium are the 
essential compounds in the Welsbach mantles. In making 
the mantles, the cotton bag is dipped into a solution of 
thorium and cerium nitrates, and then burned. The cotton 
is destroyed by heat and the nitrates are converted into a 
firm mass of oxides, which retain the shape of the mantle. 
The proportion is about 1 per cent cerium oxide (Ce0 2 ) 
and 99 per cent thorium oxide (Th0 2 ), this being the proper 
mixture for a brilliant flame of suitable color. (See also 
Bunsen Flame, Chapter XVI.) 

Thorium compounds are radioactive. (See Radioactivity.) 

Problems and Exercises 

1. What is the per cent of lead in (a) galena (PbS), (b) cerus- 
site (PbCOs), (c) anglesite (PDSO4), (d) crystallized lead acetate 
(Pb(C2H 3 02)2.3H 2 0)? 

2. How much litharge may be made from 40.5 gm. of lead? 
(Assume Pb + O = PbO.) 

3. What is the per cent of tin in (a) tinstone (Sn02), (b) stannous 
chloride (SnCl2), (c) stannic chloride (SnCU) ? 

4. If 100 gm. of tin form 127.1 gm. of stannic oxide (SnC>2), 
what is the atomic weight of tin? 

5. By analysis 100 gm. of lead monoxide yielded 7.1724 gm. of oxy- 
gen. What is the atomic weight of lead ? 

6. The formulas of cerium oxide and thorium oxide are Ce02 and 
Th0 2 ; what are the formulas of the corresponding chlorides, nitrates, 
and sulphides? 

7. Calculate the solubility product of lead chromate on the as- 
sumption that the ionization is 100 per cent and the molar solubility 
is .0000004. 



CHAPTER XXXI 
Manganese 

Occurrence. — This metal is not found free in nature, 
but its oxides are widely distributed and rather abundant. 
The chief compound is manganese dioxide (pyrolusite, 
Mn0 2 ). Other native compounds of manganese are braunite 
(Mn 2 3 ), hausmannite (Mn 3 4 ), manganite (MnO(OH) or 
Mn0 2 H), and rhodocroisite (MnC0 3 ). 

Preparation, Properties, and Uses. — Manganese is pre- 
pared by heating manganese dioxide with charcoal in an 
electric furnace, but the purest quality is made by the 
aluminothermic method. 

The equation for the latter change is — 

4A1 + 3Mn0 2 = 3 Mn + 2 A1 2 3 

Aluminium Manganese Manganese Aluminium 

Dioxide Oxide 

The metal is gray-red, hard, and brittle. If pure, it is 
permanent in the air. It interacts with dilute acids. Its 
chief property is the ready formation of alloys, especially 
with iron, copper, zinc, and nickel. It melts at 1225° C. 

Alloys of Manganese and iron are extensively used in the 
manufacture of Bessemer steel. (See Steel.) Spiegel iron 
contains from 5 to 20 per cent of manganese, while ferro- 
manganese contains 20 per cent or more. Other alloys are 
finding applications, now that pure manganese is available. 

Manganese Dioxide, Mn0 2 , is the most abundant and im- 
portant compound. It is a black solid and is often called 

517 



518 INORGANIC CHEMISTRY 

black oxide of manganese. When heated to a high tem- 
perature, it yields oxygen; and when heated with hydro- 
chloric acid, the two compounds interact, forming man- 
ganous chloride, chlorine, and water, thus : — 

Mn0 2 + 4HC1 s MnCl 2 + Cl 2 + 2H 2 

Manganese Hydrochloric Manganous Chlorine Water 

Dioxide Acid Chloride 

It colors glass and borax a beautiful amethyst, and is often 
used in glass making to neutralize the green color that 
would be caused by iron compounds in the sand. Large 
quantities are used in the manufacture of oxygen, chlorine, 
glass, and manganese alloys and compounds. 

A borax bead is colored amethyst by manganese com- 
pounds in the oxidizing flame but is colorless in the reduc- 
ing flame — a test for manganese. 

The manganese dioxide used in the manufacture of chlorine is 
recovered by the Weldon process. The impure manganous chloride 
solution from the chlorine still is treated with calcium carbonate to 
neutralize the free acid and precipitate any iron present. Lime is 
added to the clear solution of manganous chloride, and air is blown 
into the mixture. The manganous chloride is changed into man- 
ganous hydroxide (Mn(OH) 2 ), which interacts with the oxygen (of 
the air) and lime, forming chiefly calcium manganite (CaMnOs). 
After this mixture has settled, the calcium chloride is drawn off, 
and the manganese compound, which is called " Weldon mud," is 
used to generate more chlorine. 

Manganese dioxide was used by the ancients to decolorize glass, 
but its nature was misunderstood. They confused it with an iron 
oxide called magnesia stone, and the alchemists in the Middle Ages 
gave the name magnesia to this manganese dioxide. Later they 
called it magnesia nigra, or black magnesia, to distinguish it from 
magnesia alba, or white magnesia (MgO), supposing the two were 
related. Manganese was isolated in 1774, and later was given the 
specific name manganesium, which was soon shortened to manganese. 

Potassium Permanganate, KMn0 4 , is a dark purple, 
glistening, crystalline solid, though the crystals sometimes 



MANGANESE 519 

appear black, with greenish luster. It is very soluble in 
water, and solution is red, purple, or black, according to 
the concentration, Potassium permanganate gives up its 
oxygen readily and is frequently used as an oxidizing agent 
in the laboratory. It is also used as a disinfectant, a medi- 
cine, in bleaching and dyeing, in coloring wood brown, and 
in purifying gases, such as hydrogen, ammonia, and carbon 
dioxide. 

Potassium permanganate is manufactured by oxidizing a 
mixture of manganese dioxide and potassium hydroxide, 
and treating the resulting potassium manganate with sul- 
phuric acid, carbon dioxide, or chlorine. The essential re- 
actions are represented thus : — 

Mn0 2 + 2 KOH + O = K 2 Mn0 4 + H 2 

Manganese Potassium Potassium 

Dioxide Hydroxide Manganate 

3K 2 Mn0 4 + 2C0 2 = 2KMn0 4 + 2K 2 C0 3 + Mn0 2 

Potassium 
Permanganate 

The uses of potassium permanganate depend mainly upon 
its oxidizing power, that is, upon the property of liberating 
nascent oxygen readily. In an acid solution the action is 
represented thus : — 

2KMn0 4 + 3H 2 S0 4 = 50 + 2MnS0 4 + K 2 S0 4 + 3H 2 

Potassium Sulphuric Oxygen Manganese Potas- Water 

Permanganate Acid Sulphate sium Sul- 

phate 

In neutral or alkaline solutions the action is as follows : — 

2 KMn0 4 + 5 H 2 = 3 O + 2 Mn(OH) 4 + 2 KOH 

The liberated oxygen oxidizes organic matter or any other 
oxidizable substance, and the solution becomes brown or 
colorless, owing to the reduction of the potassium per- 
manganate and the transformation into manganese com- 
pounds having a faint color or none. 



520 INORGANIC CHEMISTRY 

Compounds of Manganese are numerous, often complex, 
and closely related. There are four oxides besides man- 
ganese dioxide. Three manganous compounds are impor- 
tant, the chloride (MnCl 2 ), the sulphate (MnS0 4 ), and the 
sulphide (MnS) ; in these and other manganous compounds 
the manganese plays the role of a metal. The chloride and 
sulphate are pink, crystalline salts. The sulphide is obtained 
as a flesh-colored precipitate by adding ammonium sulphide 
to the solution of a manganous salt, the color distinguish- 
ing it from all other sulphides; its formation is often used 
as a test for manganese. Manganates are salts of the hypo- 
thetical manganic acid (H 2 Mn0 4 ); the manganese in them 
acts as a non-metal. Potassium manganate (K 2 Mn0 4 ) is 
obtained as a green mass by fusing a mixture of a man- 
ganese compound, potassium hydroxide (or carbonate), and 
potassium nitrate. Its formation on a small scale consti- 
tutes a test for manganese. One equation for this chemical 
change is given on the preceding page ; another is as follows : 



Mn0 2 + 


K 2 C0 3 + 


= K 2 Mn0 4 


+ co 2 


Manganese 


Potassium 


Potassium 


Carbon 


Dioxide 


Carbonate 


Manganate 


Dioxide 



Sodium manganate (Na 2 Mn0 4 ) is used in solution as a disin- 
fectant, but sodium permanganate (NaMn0 4 ) is more effective 
and is sold in solution as " Condy's liquid." 

Miscellaneous. — Manganese has a variable valence. It is 
two in manganous compounds {e.g. MnO, Mn(OH) 2 , MnS, 
and MnS0 4 ), three in manganic compounds {e.g. Mn 2 3 ), 
four in manganese dioxide (Mn0 2 ) and in manganites, six 
in manganates {e.g. K 2 Mn0 4 ), and seven in permanganates 
{e.g. KMn0 4 ). The valence of the radical Mn0 4 is two in 
manganates and one in permanganates. 

Manganese compounds yield several kinds of ions, e.g. the 



MANGANESE 521 

delicate pink manganese ion (Mn ++ ), and the purple per- 
manganate ion (Mn0 4 "). 

The atomic weight of manganese is 54.93. Manganese 
apparently occupies an isolated position in Group VII of the 
periodic classification. 

Problems and Exercises 

1. Calculate the weight of manganese in (a) 1 metric ton of pyro- 
lusite (85 per cent pure), (b) 1 kg. of manganous chloride, and (c) 275 
gm. of potassium manganate. 

2. Write the formulas of the following: Manganese carbonate, 
manganese heptoxide, barium permanganate, manganic acid, potas- 
sium manganese alum. Calculate the per cent of manganese in 
three of these compounds. 

3. How much manganese dioxide is needed to prepare a ton 
(2000 lb.) of potassium permanganate? 

4. What (a) weight and (6) volume (standard conditions) of 
oxygen are produced by the interaction of sulphuric acid and 90 gm. 
of potassium permanganate ? 

5. What is the atomic weight of manganese, if 10.6647 gm. of 
manganous oxide (MnO) yield 22.6875 gm. of manganese sulphate 
(MnS0 4 ) ? (Use exact atomic weights.) 

6. From the following data, find the atomic weight of manganese 
and the number of atoms in the compound analyzed : An analysis 
of an oxide of manganese yielded 69.62 parts of manganese and 30.38 
parts of oxygen ; specific heat of manganese is .1217. 



CHAPTER XXXII 

Chromium — Uranium — Radioactivity 

Occurrence. — Metallic chromium is never found free. Its 
chief ore is ferrous chromite, or chrome iron ore, Fe(Cr0 2 )2. 
Native lead chromate (crocoite, or crocoisite, PbCr0 4 ) is less 
common. Traces of chromium occur in many green miner- 
als and rocks, e.g. emerald, serpentine, and verde antique 
marble. 

The name chromium comes from the Greek chroma, meaning color, 
and emphasizes the fact that most chromium compounds have con- 
spicuous colors. 

Preparation, Properties, and Uses. — Chromium was a 
rare metal until Moissan prepared it, in 1894, by heating a 
mixture of chromite and carbon in an electric furnace. The 
product contained carbon and was refined by fusing it with 
lime. Pure chromium is now prepared by the alumino- 
thermic method. (See Thermit.) 

Chromium is a lustrous gray metal. It takes a good 
polish, which is not removed by exposure to air. It is hard 
and brittle, and can be polished without difficulty. Its 
specific gravity is about 6.9. Its melting point is 1510° C. 
The heated metal burns in oxygen, forming green chromic 
oxide (Cr 2 3 ). 

Chromium is used to harden the steel that is made into 
armor plates, hard tools, projectiles, safes and vaults, and 
certain parts of machines used to crush gold-bearing quartz. 
This hardened steel is called chromium steel. The com- 
mercial form of chromium is usually an alloy of 65 to 80 

522 



CHROMIUM 523 

per cent chromium, a little carbon, and the rest iron; this 
alloy is called ferrochrome. 

Compounds of Chromium are numerous, some are com- 
plex, many pass readily into one another, and a few have 
industrial applications. The most important are potassium 
chromate, potassium dichromate, chrome alum, and lead 
chromate. 

Potassium Chromate, K 2 Cr0 4 , and Potassium Dichromate, 
or Bichromate, K 2 Cr 2 7 . — These compounds are manu- 
factured from chrome iron ore. The crushed ore is mixed 
with lime and potassium carbonate and roasted in a rever- 
beratory furnace; air is freely admitted and the mass fre- 
quently raked. By this operation the ore is oxidized into a 
mixture of calcium and potassium chromates. The mass is 
cooled, pulverized, and treated with a hot solution of potas- 
sium sulphate, which changes the calcium chromate into 
potassium chromate, thus : — 

CaCr0 4 + K 2 S0 4 = K 2 O0 4 + CaS0 4 

Calcium Potassium Potassium Calcium 

Chromate Sulphate Chromate Sulphate 

The solution of potassium chromate is filtered, concentrated, 
and changed by sulphuric acid into potassium dichromate; 
the latter is purified by recrystallization from w r ater. Potas- 
sium chromate is a lemon-yellow, crystalline solid, which 
contains no water of crystallization. It is very soluble in 
water and gives a yellow solution. Acids change it into the 
dichromate, thus : — 

2K 2 Cr0 4 + H 2 S0 4 = K 2 Cr 2 7 + K 2 S0 4 + H 2 

Potassium Sulphuric Potassium Potassium Water 

Chromate Acid Dichromate Sulphate 

Potassium chromate is also formed as a yellow mass by 
fusing intensely on porcelain (or platinum) a mixture of a 



524 INORGANIC CHEMISTRY 

chromium compound, potassium carbonate, and potassium 
nitrate. When the mass is boiled with an excess of acetic 
acid to decompose the carbonate and expel the carbon 
dioxide, and then added to a lead salt solution, yellow lead 
chromate is precipitated. This experiment is often used as 
a test for chromium. (See also Lead Chromate.) Potassium 
dichromate is a red solid, which is prepared as described 
above. It is less soluble in water than potassium chromate, 
and yields a pale yellow or red solution according to the 
concentration. Hydroxides change it into the chromate, 
thus : — 

K 2 Cr 2 7 + 2KOH = 2K 2 Cr0 4 + H 2 

Potassium Potassium Potassium Water 

Dichromate Hydroxide Chromate 

Potassium dichromate is used in dyeing, calico printing, 
and tanning, in bleaching oils, and in manufacturing chro- 
mium compounds and dyestuffs. Its uses depend mainly 
upon the fact that it is an oxidizing agent. When potas- 
sium dichromate and sulphuric acid are mixed, the equation 
for the reaction may be written : — 

K 2 Cr 2 7 + 4H 2 S0 4 = 3 + K 2 S0 4 + Cr 2 (S0 4 ) 3 + 4H 2 

Potassium Sulphuric Oxygen Potassium Chromium 

Dichromate Acid Sulphate Sulphate 

Some oxidizable substance, however, such as ferrous sul- 
phate, must be present to use up the liberated oxygen, 
thus : — 

6FeS0 4 + 3H 2 S0 4 + 30 = 3Fe 2 (S0 4 ) 3 + 3H 2 

Ferrous Ferric 

Sulphate Sulphate 

The complete chemical change is often expressed as follows: — 

K 2 Cr 2 7 + 7 H 2 S0 4 + 6 FeS0 4 = 3 Fe 2 (S0 4 ) 3 + K 2 S0 4 

+ Cr 2 (S0 4 ) 3 + 7H 2 



CHROMIUM 525 

Potassium chromate and dichromate are being replaced somewhat 
by the corresponding sodium salts, because the latter are cheaper, 
more soluble, and have analogous properties. The potassium salts are 
anhydrous, but crystallized sodium chromate is Na 2 Cr0 4 . 10 H 2 
and crystallized sodium dichromate is Na 2 Cr20 7 • 2 H2O. 

Chrome Alum, K 2 Cr 2 (S0 4 ) 4 . 24 H 2 0, is a purple, crystal- 
line solid. It is analogous in composition and similar in 
properties to ordinary alum, but it contains chromium in- 
stead of aluminium. It can be prepared by mixing potas- 
sium and chromium sulphates in the proper proportion, or 
by passing sulphur dioxide into a solution of potassium 
dichromate containing sulphuric acid. The commercial sub- 
stance is a by-product obtained in the manufacture of the 
dyestuff alizarin. Chrome alum is used as a mordant in 
dyeing and calico printing and in tanning. 

Lead Chromate, PbCr0 4 , is a bright yellow solid, formed 
by adding potassium chromate or dichromate to a solution 
of a lead salt. An equation for the chemical change is — 

K 2 Cr 2 7 + 2 Pb(N0 3 ) 2 + H 2 = 2 PbCr0 4 + 2 KN0 3 + 2 HN0 3 

Potassium Lead Water Lead Potassium Nitric 

Dichromate Nitrate Chromate Nitrate Acid 

It is known as chrome yellow, and is used in making yellow 
paint. When boiled with sodium hydroxide, lead chromate 
is changed into a basic chromate (PbCr0 4 . PbO . H 2 0) called 
chrome red or chrome orange, depending on the color. 

The precipitation of lead chromate by the interaction of 
a dissolved lead salt and a dissolved chromate (or dichro- 
mate) is often used as a test for chromium. 

Chromium forms Four Series of Compounds, the chro- 
mous, the chromic, the chromites, and the chromates 
(mono- and di-). In the chromous and chromic compounds, 
chromium acts as a metal, but in chromites and chromates 
it acts as a non-metal. 



526 INORGANIC CHEMISTRY 

Chromous Compounds may be regarded as derived from 
chromous oxide (CrO). As a class they are so easily oxidized 
into chromic compounds that they are difficult to prepare 
and keep. 

Chromic Compounds may be regarded as derivatives of 
chromic oxide (Cr 2 3 ). This is a bright green powder pre- 
pared by heating chromic hydroxide (Cr(OH) 3 ), and is the 
basis of the chrome green pigments used to color glass 
and ornament porcelain. When chromium compounds are 
heated with borax, they color the bead green in both flames, 
owing to the formation of this oxide (Cr 2 3 ). If potassium 
dichromate and boric acid are mixed and heated, and then 
treated with water, a hydrated chromic oxide is formed 
called Guignet's green (Cr 2 3 . 2 H 2 0) ; it gives a permanent 
color and is extensively used. There are several chromic 
hydroxides. The typical one has the composition repre- 
sented by the formula Cr(OH) 3 . It is a bluish solid formed 
by the interaction of a chromic compound (e.g. chrome 
alum) and an alkaline hydroxide or sulphide. The precipi- 
tation of chromic hydroxide by ammonium sulphide is due 
to the fact that chromic sulphide (which we might expect 
to be formed) hydrolyzes, yielding chromic hydroxide and 
hydrogen sulphide. (Compare Aluminium Hydroxide, Chap- 
ter XXIX.) 

Chromic hydroxide is soluble in an excess of sodium 
(or potassium) hydroxide. That is, it is changed into 
a soluble chromite, just as aluminium hydroxide forms 
soluble aluminates. Unlike aluminates, however, the chro- 
mites are changed back into chromic hydroxide by boiling. 
Other chromic salts are chromic chloride (CrCl 3 ), chromic 
sulphate (Cr 2 (S0 4 ) 3 ), and potassium chromium sulphate or 
chrome alum (K 2 Cr 2 (S0 4 ) 4 . 24 H 2 0). The valence of chro- 
mium in chromic compounds is three. 



URANIUM 527 

Chromites may be regarded as salts of an acid having the 
composition corresponding to HCr0 2 ; native chromite 
(FeCr 2 4 or Fe(Cr0 2 ) 2 ) is an iron salt of this acid. 

Chromates and dichromates start theoretically from chro- 
mium tri oxide (Cr0 3 ). This is the anhydride of the hypo- 
thetical chromic acid (H 2 Cr0 4 ). When concentrated sul- 
phuric acid is added to a saturated solution of potassium 
dichromate (or chromate), chromium tri oxide (Cr0 3 ) sepa- 
rates as long, bright red crystals, thus: — 

K 2 Cr 2 7 + H 2 S0 4 = 2 Cr0 3 + K 2 S0 4 + H 2 

Potassium Sulphuric Chromium Potassium Water 

Dichromate Acid Trioxide Sulphate 

It is sometimes called chromic acid, and is a vigorous 
oxidizing agent. The valence of chromium in chromates 
and dichromates is six; the radicals Cr0 4 and Cr 2 7 have 
the valence two. 

Miscellaneous. — Chromium compounds yield several kinds 
of ions, e.g. the violet chromic ion (Cr + + + ), the yellow 
chromate ion (Cr0 4 ~~), and the red dichromate ion 

(CrA-). 

Molybdenum (Mo), Tungsten (W), and Uranium (U) are rare metal- 
lic elements related to chromium. Most of their compounds have 
only scientific interest, though some have analytical or industrial uses. 
Ammonium molybdate ((XH 4 ) 2 Mo0 4 ) is used in the laboratory to de- 
termine the amount of phosphorus in fertilizers and iron. Tungsten 
is used to harden steel, and as the filament of electric light bulbs; 
sodium tungstate (Na 2 W0 4 ) finds application in rendering cloth fire- 
proof. Uranium compounds are obtained from uraninite and pitch- 
blende, the latter mineral now being the chief source. The element 
forms many compounds and the important ones are the oxides, sodium 
uranate (Na 2 U 2 7 . 6 H 2 0), and uranyl nitrate (U0 2 (N0 3 ) 2 . 6 H 2 0); 
from the nitrate other salts are prepared. Sodium uranate is some- 
times called uranium yellow, and it is used to make uranium glass. 
Such glass is green by transmitted light and yellow by reflected light. 
Uranium is a radioactive element.. 



528 INORGANIC CHEMISTRY 



Problems and Exercises 

1. What is the per cent of chromium in (a) lead chromate 
(PbCr0 4 ), (b) chrome ironstone (Fe(Cr0 2 )2), (c) chromic oxide 
(Cr 2 3 )? 

2. How many grams of lead chromate can be made from 438 gm. 
of potassium dichromate ? 

3. What is the solubility product of barium chromate if the 
ionization is 89 per cent and the molar solubility is .000014? 



Radioactivity 

Historical. — It was found about 1896 that uranium com- 
pounds affect a photographic plate and discharge an electro- 
scope. Numerous experiments by Becquerel, Curie, Ruther- 
ford, and others showed that these effects were probably due 
to radiations like X-rays and Rontgen rays. Furthermore, 
experiments indicated that this power of radiation belongs 
to the uranium itself, and that the radiations are emitted 
spontaneously without the aid of any outside agency. Other 
substances, subsequently shown to be radium compounds, 
were found to possess similar properties, and, as stated in 
the preliminary discussion of radium, such substances are 
said to exhibit radioactivity. Soon after (1898) an exten- 
sive examination of many minerals showed that certain min- 
erals containing uranium, especially pitchblende, are more 
radioactive than uranium itself. A radioactive substance 
was extracted from pitchblende by M. and Mme. Curie, and 
the elementary constituent in it was named radium. 

Interpretation of Radioactivity. — Radioactivity is not due, 
as was first supposed, to radiations, but to the spontaneous 
emission by radioactive substances of two kinds of particles, 
which are called alpha (a) and beta (/?) ; the emission of beta 
particles is accompanied by pulsations in the ether known as 
gamma (y) rays, 




MADAME CURIE 



RADIOACTIVITY 529 

The alpha particles are shot off in a stream which moves 
with a velocity averaging about 18,000 miles a second. The 
path of single alpha particles in a special apparatus has been 
photographed. Alpha particles bear a positive charge of 
electricity and can be readily detected by a delicate electro- 
scope. Many of the electrical phenomena of radioactive 
substances are due to alpha particles. Alpha particles are 
four times as heavy as hydrogen atoms, and they are iden- 
tical with charged atoms of the element helium. 

An instrument called the spinthariscope shows vividly that alpha 
particles are being shot off continuously by a radium compound. It 
is a small microscope with a screen and pin opposite the lens ; the 
screen is coated with zinc sulphide and on the needle there is a min- 
ute quantity of radium bromide. Upon looking through the lens, 
minute flashes of light are seen on the screen. They are due to the 
alpha particles which produce fluorescence in the zinc sulphide. 

The beta particles consist of a stream of electrons, i.e. par- 
ticles of negative electricity, moving with a varying velocity 
which is sometimes nearly as great as the velocity of light 
(186,000 miles a second). Beta particles are very light, their 
weight being about T gVo" °f the weight of a hydrogen atom. 
To the beta particles are ascribed most of the photographic 
effects of radioactive substances. 

The gamma rays are not material particles, but like X-rays 
are pulsations in the ether. The curative effect of radium is 
believed to be due to gamma rays. 

Alpha particles move in straight lines and penetrate air 
to a depth of 3 to 8 centimeters. They are almost entirely 
stopped by a thin sheet of paper and by aluminium leaf 
.1 millimeter thick. Beta particles move in straight lines at 
first, but soon in curved lines, owing to collisions with the 
relatively heavier molecules of the gases of the air. They 
penetrate air to a less depth than alpha particles ; they pass 
through gold leaf but are stopped by aluminium 1 centimeter 



530 INORGANIC CHEMISTRY 

thick. The gamma rays are the most penetrating. They 
pass readily through thick layers of metals ; glass tubes con- 
taining radium salts are enclosed in lead vessels to absorb 
the gamma rays. The stoppage of so many and such 
rapidly moving particles by the air, metals, and the radio- 
active substance itself develops heat. It is estimated that 
one gram of radium (metal, in combination) would produce 
spontaneously about 120 calories per hour. 

Disintegration of Radioactive Compounds. — Uranium and 
radium are like most of the other chemical elements in general 
respects, but they differ in one conspicuous property, viz. 
atomic instability. That is, the atoms of uranium, radium, 
and a few other elements are spontaneously disintegrating. 
Uranium is the parent, so to speak, of a series of products 
formed step by step by disintegration. Among these prod- 
ucts are the elements radium and niton. The rate of dis- 
integration varies widely and is not affected by conditions, 
that is, it is spontaneous. In a unit time a definite fraction 
of a product disintegrates ; it is customary in describing dis- 
integration to state the time in which half the amount would 
disintegrate. Thus, the half period of radium is 2400 years 
and of niton is 5.55 days. 

Certain interesting facts bearing on disintegration are well 
established. For example, uranium ores contain amounts of 
radium that are proportional to the uranium. Again, a ura- 
nium compound gradually recovers its whole radioactivity 
after the radium has been removed. Furthermore, helium 
is produced at different steps in the disintegrating process; 
the helium given off by radium compounds has been collected, 
studied, and its rate of production measured (164 cubic 
millimeters per gram of radium (metal) a year). 

The products formed by successive disintegration in the 
uranium series include electrons and helium besides radium, 



RADIOACTIVTY 531 

niton, and certain less known elements. In other words, the 
product of the disintegration of an atom of an element may 
be an atom of another element or an atom plus either an elec- 
tron or a charged helium atom. The uranium (Uranium 
— 1) series is as follows : — 

Uranium -1 (238) — ^ Uranium -X x (234) — >- 
Uranium - X 2 (234) — >- Uranium - 2 (234) — ^ 

Ionium (230) — >- Radium (226) — >- 

Niton (222) — ^Radium -A (218) — >- 

Radium-B (214) — ^Radium-C (214) — >- 

Radium -d (214) — >- Radium -D (210) — >- 

Radium-E (210) — ^Radium-F (210) — >- 
Radium -G (206). 

In the above series it will be noticed that the atomic weight 
(given in the parenthesis) of certain elements is 4 less than 
the atomic weight of the preceding element, e.g. niton (222) 
and radium (226). In such cases the atomic disintegration is 
accompanied by the expulsion of a charged helium atom (i.e. 
an alpha particle) having the atomic weight 4. On the other 
hand, the atomic weight of certain elements is the same, e.g., 
radium — D, — E, — F (210). In these cases an electron (i.e., 
a beta particle) is expelled. Although certain elements: 
have the same atomic weight, they are distinct chemical ele- 
ments and form analogous compounds which have different 
chemical properties. Moreover, certain elements in this 
series have chemical properties identical with those of ele- 
ments already known ; it is believed that radium — G is the 
final product of disintegration. 

Other Radioactive Elements. — Thorium is radioactive, 
though to a much less degree than radium. It furnishes a 
series of disintegration products similar to uranium. Ac- 
tinium and polonium are also radioactive elements. 



CHAPTER XXXIII 

Iron, Nickel, and Cobalt 

Iron 

Iron is the most useful of all metals. It has been known 
for ages, though not so long as the other common metals, 
and has been indispensable in the development of the human 
race. 

The symbol of iron, Fe, is from the Latin word ferrum. From 
ferrum are derived the words ferric and ferrous, which give the 
corresponding forms ferri- and ferro- (found in such words as ferri- 
cyanide, ferrocyanide, etc.). 

Occurrence of Iron. — Uncombined iron is found only in 
meteorites, which fall upon the earth from remote regions 
in space, and in certain volcanic rocks. Combined iron is 
abundant and widely distributed, constituting about 4.5 
per cent of the earth's crust. It is found in most rocks 
and many minerals, in the soil, in springs and natural 
waters, in chlorophyll (the green coloring matter of plants), 
and in haemoglobin (the red coloring matter of the blood). 
The chief ores of iron are hematite (Fe 2 3 ), limonite 
(2Fe 2 3 .3H 2 0), magnetite (Fe 3 4 ), and siderite (FeC0 3 ). 
These ores often contain some impurity, such as silica, clay, 
calcium or magnesium carbonate, and small quantities of 
compounds of sulphur, phosphorus, and manganese. 

Other abundant native compounds of iron are pyrites 
(FeS 2 ), pyrrhotite (varying from Fe 6 S 7 to Fe n S 12 ), and the 
copper-iron sulphides (chalcopyrite, CuFeS 2 , and bornite, 
Cu 3 FeS 3 ). They are not used to any extent as a source of 
iron. 

532 



IRON 



533 



The United States leads the world in the production of iron ore, 
the annual output for the last few years being about 50,000,000 tons. 
This vast quantity conies from twenty-five different states, but the 
bulk is mined in Minnesota, Michigan, Alabama, Wisconsin, Ten- 
nessee, Virginia, West Virginia, and Colorado. The most abundant 
ore is the red hematite, which comes chiefly from the Lake Superior 
region (Fig. 79); large quantities are mined in Alabama and Ten- 
nessee. The latter states, together with Virginia and West Virginia, 




Fig. 79. — Deposits of iron and copper near Lake Superior. The iron 
regions, known as ranges, are Marquette (1), Menominee (2), Gogebic (3), 
Vermilion (5), Mesabi (6). No. 4 is the copper region. 

furnish most of the limonite or brown iron ore. Pennsylvania, New 
Jersey, and New York contribute most of the magnetite, though 
some is mined also in Michigan. The carbonate ores, which con- 
stitute less than 1 per cent of the output, come mainly from Ohio, 
Maryland, and New York. Improvements in the machinery and 
methods used in mining and transporting iron ore have reduced its 
cost and facilitated its production. Thus, at an incredibly small 
expense, ore from the Lake Superior region is raised from open pits 
by steam shovels, dumped into large ears, carried to shipping ports 
on the lakes, dumped again into huge bunkers, dropped down chutes 
into big freight steamers (many of which hold 6000 tons), which 
carry it to southern ports on Lake Michigan, though the large part 



534 



INORGANIC CHEMISTRY 



is sent to ports on the south shore of Lake Erie and forwarded by 
rail to Pittsburg, Pennsylvania. This city is the center of iron and 

steel industries. Birmingham, 
Alabama, is the center of the 
industry in the South, because 
near it the necessary ore, coal, 
and limestone are conveniently 
located. 

Metallurgy of Iron. — 

Iron is extracted most 
easily from its oxides. 
Therefore the ore, unless 
it is very pure hematite, is 
first crushed and roasted to 
change it into ferric oxide 
(Fe 2 3 ) as far as possible. 
The ore is then mixed 
with coke and limestone 
or sand, and smelted in a 
blast furnace. The carbon 
reduces the oxide to metal- 
lic iron, which collects as a 
liquid at the bottom of the 
furnace beneath the slag 
formed by the limestone 
and impurities. The blast 
furnace (Fig. 80) is a huge 
circular tower, from forty 
to ninety feet high, and 
about thirteen feet in di- 
ameter at the largest part. 
It is built of iron and lined 
with fire brick. Pipes near 
the bottom, called tuyeres, 




Fig. 80. — Blast furnace : A, throat; B, 
bosh ; C, crucible where the melted iron 
collects; D, pipes for hot-air blast; E, 
escape pipe for gases which do not 
escape through the "down comer"; 
G, cup; H, cone; N, trough for draw- 
ing off slag; T, tuyere; I, hole through 
which iron is withdrawn. 



allow large quantities of hot air to be forced into the furnace 



IRON 535 

up through the contents, thereby producing the high tem- 
perature required in the smelting; while another pipe at the 
top not only permits the escape of the hot gaseous products, 
but conducts them into a series of pipes which lead to dif- 
ferent parts of the plant, where the hot gases are utilized to 
heat the air which is blown through the furnace. The blast 
pipes correspond to the bellows used by the blacksmith, 
and the exit pipe to the chimney, except that gases escaping 
through chimneys are usually wasted. 

When the blast furnace has been heated to the proper 
temperature, or is already in operation, the mixture or charge 
is carried to the top by machinery and introduced into the 
furnace by dumping it upon the cone-shaped cover; the 
weight lowers the cover, which flies back tightly into place 
after the materials roll into the furnace. The charge con- 
sists of the proper mixture of ore, fuel, and flux. The ore, 
as stated above, is usually hematite. The fuel is coke, 
or coke mixed with coal. The flux varies with the impurities 
in the ore; thus, it is limestone if the ore contains silica or 
clay, but sand if the impurities are calcium or magnesium 
compounds. It is usually limestone. The object of the 
flux is to remove the impurities (just mentioned) from the 
charge in the form of readily fusible silicates called slag. 
As the smelting proceeds, the contents of the furnace slowly 
descend and are changed into gases, iron, and slag. The 
gases rise through the mass and escape by pipes, the solids 
become pasty at first and then liquid, the iron finally drop- 
ping through the slag into the crucible at the bottom of the 
furnace, where both are tapped off through separate open- 
ings. Fresh charges of definite weight and proportions 
are added at regular intervals, and the whole operation 
continues without interruption for months or even years. 

The iron from the furnace is usually poured into molds 
of sand or iron and allowed to solidify. Such iron is called 



536 INORGANIC CHEMISTRY 

pig iron or cast iron. In some plants the molten iron is 
run directly from the blast furnace into huge vessels called 
converters and made at once into steel (see below) ; in other 
plants it is kept molten in huge tanks until needed. 

The chemical changes involved in the metallurgy of iron 
are numerous. In general, the iron oxide is reduced to 
metallic iron largely by carbon monoxide. The carbon of 
the fuel at first forms carbon dioxide with the oxygen of 
the air blast. But the dioxide is soon reduced by the hot 
carbon to the monoxide, which interacts with the ore, thus : — 

3Fe 2 3 + CO = 2Fe 3 4 + C0 2 

Ferric Carbon Ferrous-ferric Carbon 

Oxide Monoxide Oxide Dioxide 

Fe 3 0< + CO = 3FeO + C0 2 

Ferrous 
Oxide 

FeO + CO = Fe 4- CO, 

At this stage the ore, though not wholly reduced, becomes 
soft and porous, and as the mass sinks into the hottest part 
of the furnace the reduction is completed, thus : — 

FeO + C = Fe 4- CO 

The iron now combines with a small percentage of carbon, 
melts, and sinks through the molten slag. This iron contains 
small amounts of carbon, silicon, manganese, phosphorus, 
and sulphur. 

Varieties of Iron. — The iron we use and speak of is not 
pure iron, but a mixture or compound of iron with other 
elements, chiefly carbon. It is customary to speak of three 
varieties of iron, — cast iron, wrought iron, and steel. 
This classification is based chemically upon the per cent of 
carbon they contain, though their physical properties are 
also modified by the presence of silicon, phosphorus, sulphur, 



IRON 537 

and manganese, as well as by the method of manufacture. 
The different varieties are closely related, and pass easily 
and gradually into each other. Commercially, there are 
several kinds of cast iron and many kinds of steel. 

Cast Iron is the most impure variety. It contains, besides 
carbon, the impurities already mentioned. The carbon 
varies from 3 to 5 per cent, the silicon and manganese are 
each about 3 per cent, while the proportion of phosphorus and 
sulphur is small. If molten iron is cooled suddenly, the prod- 
uct, which is very brittle, is called white cast iron ; the car- 
bon in it is mostly in the form of a carbide (cementite, Fe 3 C). 
By cooling slowly, much of the carbon remains uncombined 
as hard crystals known as graphite carbon, and the color of 
the iron is gray; this kind is gray cast iron. It is softer 
than the white variety, and melts at a lower temperature. 
Although cast iron is brittle, it will withstand great pressure. 
Owing to its crystalline structure, it cannot be welded or 
forged; that is, hot pieces cannot be united, nor be shaped 
by hammering. But it is extensively used to make castings. 
This is the kind of iron used in an ordinary iron foundry. 
The iron, which melts at about 1200° C. (depending upon the 
impurities), is heated in a furnace similar to a blast furnace, 
and when molten is poured into sand molds of the desired 
shape. Stoves, pipes, pillars, railings, parts of machines, 
and many other useful objects are made of cast iron. 

Cast iron containing 5 to 20 per cent of manganese is called 
spiegel iron, while ferro-manganese contains from 20 to 
85 per cent of manganese. Both are used in making steel. 

Wrought Iron is the purest variety of commercial iron. 
It contains not more than 0.5 per cent of carbon and some- 
times only 0.06 per cent, the average being 0.15 per cent. 
It is tough, malleable, and fibrous, and can be bent. Unlike 
cast iron, it does not withstand pressure, but it will sustain 



538 INORGANIC CHEMISTRY 

great weight. An iron wire will sustain the weight of 
nearly a mile of itself. Wrought iron melts at such a high 
temperature (1550° to 2000° C.) that it is not used for 
casting; it softens at a relatively low temperature (about 
1000° C), can be forged and welded, and is often called 
malleable iron. It may be seen undergoing these opera- 
tions in a blacksmith's shop. It can also be rolled into 
sheets and plates and drawn into fine wire; in these forms 
the metal is very tough. Wrought iron is made into wire, 
sheets, rods, nails, spikes, bolts, chains, anchors, horse- 
shoes, tires, and agricultural implements. It is less im- 
portant than formerly, since it is being largely replaced by 
steel. 

Wrought iron is made from cast iron by burning out most 
of the impurity. Cast iron together with a little scrap iron 
and flux is melted in a furnace, much like a reverberatory 
furnace, lined on the bottom and sides with iron ore (ferric 
oxide, Fe 2 3 ). The impurities are oxidized partly by the 
oxygen of the air, but mainly by the oxygen of the iron oxide ; 
the silicon, phosphorus, sulphur, and manganese pass off 
in the slag ; most of the carbon, escapes as an oxide, though 
a small amount remains in the iron. As these elements are 
removed, the mass becomes pasty, owing to the higher 
melting point of the pure iron. It is now stirred vigorously, 
or " puddled. " At the proper time lumps called "blooms" 
are removed and hammered, or more often rolled between 
ponderous rollers. This operation removes the slag, and if 
the rolling is repeated, the quality of the iron is improved; 
the final rolling often leaves the iron in the shape desired 
for market. 

Steel is usually intermediate between cast iron and wrought 
iron as far as its proportion of carbon is concerned. Many 
grades of steel are manufactured, and their physical prop- 



IRON . 539 

erties do not depend merely upon the presence of a small 
proportion of carbon and other elements, especially phos- 
phorus, silicon, and certain metals, but to a considerable 
extent upon the method of manufacture and subsequent 
treatment. Considered from the standpoint of composition, 
it may be said that in general the higher the percentage of 
carbon, the harder the steel. In soft or mild steel the carbon 
is seldom more than .2 per cent, while in hard steel the 
amount may be as high as 1.5 per cent. 

Manufacture of Steel. — The aim in the manufacture of 
steel is to prepare a product containing little or no sulphur, 
phosphorus, and silicon, but the desired proportion of carbon. 
This is accomplished by several processes, viz. the Bessemer, 
open-hearth, crucible, and cementation. 

(1) The Bessemer process, which is quite generally used, 
was devised about 1860, and has practically revolutionized 
steel making. The process consists in burning out the 
impurities in cast iron by forcing air through the molten 
metal, and then adding just enough iron of known com- 
position to give the desired proportion of carbon. The 
operation is carried on in a converter (Fig. 81). This is 
a huge, pear-shaped vessel, supported so that it can be 
tipped into different positions; it is also provided with 
holes (C, C, C) at the bottom, through which a powerful blast 
of air can be blown. It is made of thick wrought-iron plates, 
and is lined with an infusible mixture, usually rich in silica. 
The converter when in use is swung into a horizontal position 
(A) , and five to twenty tons of molten pig iron are poured 
in, often directly from the blast furnace. The air blast is 
turned on and the converter is swung back to a vertical 
position (5). As the air is forced in fine jets through the 
molten metal, the temperature rises, and the carbon, silicon, 
and manganese are oxidized. The carbon forms carbon mou- 



540 



INORGANIC CHEMISTRY 



oxide, which burns at the mouth of the converter, while the 
other oxides pass into the slag. This oxidation generates 
enough heat to keep the metal melted, and no fuel need be 
used. As soon as the impurities have been burned out, 
sufficient spiegel iron or ferromanganese is added to furnish 
the proper amount of carbon and manganese. By adding 




Air 




B 
Fig. 81. — Converter. 



spiegel iron of known composition, Bessemer steel of any- 
desired grade is produced. After the completion of the 
operation, which takes about twenty minutes, the metal is 
poured from the converter into molds to cool. These cold 
blocks of steel are called ingots. 

In the Bessemer process, as described in the preceding 
paragraph, sulphur and phosphorus are not removed from 
the steel. Both are objectionable impurities; sulphur makes 
steel brittle when hot, and phosphorus makes it brittle when 
cold. The Thomas-Gilchrist process (or basic process) is 
a modification of the Bessemer process by which the sulphur 
and phosphorus can be removed in the slag. The con- 
verter in this process is lined with burned dolomite (i.e. 
practically a mixture of lime and magnesia), called a basic 
lining ; lime is also added to the charge of pig iron, and the 



IRON 



541 



air blast is continued a little longer than in the Bessemer 
process; otherwise the operations are the same. The car- 
bon passes off as usual, and the oxides of phosphorus, silicon, 
and sulphur combine with the basic constituents of the lining 
to form a slag. This lining, after use, is known as Thomas 
slag; it is utilized as a fertilizer on account of its phosphorus 
content. 

(2) In the Siemens-Martin or open-hearth process cast 
iron, or a mixture of cast iron, scrap iron, and iron ore, in 




Fig. 82. — Open-hearth furnace. 

proper proportions, is melted by hot gases in a special 
kind of furnace called an open-hearth furnace (Fig. 82). 
The receptacle, or hearth (77), in which the charge is melted, 
is lined with sand in the acid process or with lime and magne- 
sia in the basic process, thereby permitting the removal of 
all the impurities. At the base of the furnace are duplicate 
sets of checkerwork (A, B and C, D). As the hot gases 
pass through A, B to the chimney, they heat the checker- 
work. The fuel gas is then passed through B and air 
through A. The mixture of air and gas burns and pro- 



542 INORGANIC CHEMISTRY 

duces a much higher temperature on the hearth than if 
the gaseous mixture were cool. The oxidizing flame passes 
over the charge on the hearth (H), oxidizing some of the 
impurities and keeping the mass at such a temperature 
that other impurities form a slag with the lining. Mean- 
while the hot products of combustion and the unused gases 
are passed through the checkerwork C, D and heat it. 
The fuel gas and air are then made to pass through this 
checkerwork to the hearth and out over the other checker- 
work (A, B) to the chimney. Thus the process is alternated, 
one checkerwork being cooled as the other is heated, and 
vice versa. It is only by this regenerative process, as it is 
called, that enough heat is obtained to keep the charge 
melted as it becomes purer and purer. The charge is heated 
from eight to ten hours ; when a test shows that the metal 
has the desired composition, certain metals or ferro-alloys 
are added, and the steel is quickly poured into molds and 
allowed to cool into ingots. Subsequently the ingots of this 
kind of steel (as well as of other kinds) are reheated, rolled, 
and cut into lengths of the proper size and shape, being 
then known as billets. The open-hearth process requires 
a special furnace and gas plant, and is more expensive than 
the Bessemer process, since it takes much longer. But it 
is easily controlled, and yields a tough, elastic steel, which 
is excellent for bridges, large machines, large guns, and gun 
carriages. The production of the open-hearth steel has 
increased rapidly in the last few years. 

(3) In the crucible process wrought iron is melted with 
charcoal in graphite or clay crucibles. During the melting, 
which lasts three or four hours, the iron is slowly changed 
into steel by absorbing the proper proportion of carbon. 

(4) In the cementation process wrought iron and carbon 
are heated in fire-brick boxes for several days. The trans- 
formation is the same as in the crucible process. 



IRON 543 

The four processes of making steel, just described, provide 
many grades which permit its use for countless purposes. 
In recent years special steels have been made by adding 
to steel small quantities of other metals, such as nickel, 
chromium, molybdenum, tungsten, vanadium, and manganese. 
Such steels differ somewhat in special properties, but all are 
characterized by extreme hardness, toughness, and strength. 
The metals are sometimes added directly to the molten steel, 
but more often in the state of an alloy of iron. These 
alloys contain varying but known percentages of the con- 
stituents, and are called ferrochrome, ferrosilicon, ferrotung- 
sten, etc. 

The important properties of steel are numerous. It is 
fusible and malleable, and can be forged, welded, and cast. 
It is harder, stronger, and more durable than pure iron, and 
is therefore more serviceable. But its most valuable prop- 
erty is the varying hardness which it may be made to 
acquire. If steel is heated very hot and then suddenly cooled 
by immersion in cold water or oil, it becomes brittle and very 
hard. But if heated and cooled slowly, it becomes soft, 
tough, and elastic. All grades of hardness may be obtained 
between these extremes. And if the hardened steel is re- 
heated to a definite temperature, determined approximately 
by the color the oxidized metal assumes, and then properly 
cooled, a definite degree of hardness and elasticity is ob- 
tained. This last operation is called tempering. 

Uses of Steel. — Steel is now used instead of iron for many 
purposes. High buildings, bridges, rails, cars, locomotives, 
battleships, electrical machinery, boilers, agricultural im- 
plements, wire nails, rods, hoops, tin plates, and castings 
of all kinds consume vast amounts of Bessemer and open- 
hearth steel. Crucible, cementation, and special steels are 
used in making springs, tools, cutlery, pens, and needles. 



544 INORGANIC CHEMISTRY 

Properties of Pure Iron. — Chemically pure iron can be 
obtained as a black powder by reducing the oxide with 
hydrogen. Recently very pure massive iron has been pre- 
pared electrolytically from a solution of ferrous and am- 
monium sulphates. The purest commercial form is the 
wrought iron used for piano wire. Pure iron is a silvery 
white, lustrous metal. It is softer than ordinary iron, but 
melts at a higher temperature (about 1520° C). The 
specific gravity is about 7.8. It is attracted by a magnet, 
but soon loses its own magnetism. Dry air has no effect 
upon iron, but moist air rusts it. The rusting of iron is a 
complex process. Interpreted by the electrolytic dissociation 
theory the iron first interacts with water and goes into solu- 
tion as ferrous ions (Fe ++ ), while the hydrogen ions (H + ) be- 
come hydrogen molecules and escape ; the ferrous ions com- 
bine with the hydroxyl ions (OH") left in the water and form 
ferrous hydroxide (Fe(OH) 2 ), which subsequently becomes 
iron rust. Rusting proceeds rapidly, because the film of rust 
is not compact enough to protect the metal. Iron readily 
interacts with dilute acids, and as a rule hydrogen and ferrous 
compounds are the products. 

With nitric acid various products result, according to the con- 
ditions, — ferrous nitrate and ammonium nitrate if the acid is 
cold, but ferric nitrate and oxides of nitrogen if the acid is warm. 
If a clean iron wire is dipped into fuming nitric acid and then into 
ordinary nitric acid, no action is apparent. The iron is said to be 
passive. This peculiar fact has not been adequately explained. 
Steam and hot iron interact thus : — 



3Fe + 4H 2 r 


Fe 3 4 + 4H 2 


Iron Water 


Iron Oxide Hydrogen 



The film of the oxide (FesO^, unlike iron rust, adheres firmly and 
protects the iron from further oxidation. 

Compounds of Iron. — Iron forms two important series 
of compounds, — ferrous and ferric. They are analogous 



IRON 545 

to cuprous and cupric, mercurous and mercuric compounds. 
The valence of iron is two in ferrous compounds and three 
in ferric. Ferrous compounds in an acid solution pass into 
the corresponding ferric compound by the action of oxidiz- 
ing agents, e.g. oxygen, nitric acid, potassium chlorate, 
potassium permanganate, and chlorine. Conversely, ferric 
compounds are reduced to the ferrous by reducing agents, 
e.g. hydrogen, hydrogen sulphide, sulphur dioxide, and 
stannous chloride. The passage from one series to the other 
occurs easily, especially from ferrous to ferric. The oxida- 
tion and reduction of iron compounds illustrates typically 
the broad use of these terms. Oxygen is not necessarily 
involved. Thus, ferrous and ferric chlorides pass readily 
into each other by the addition or removal of chlorine. The 
two processes are general and mutual, and may be sum- 
marized as follows: (a) Oxidation is the chemical addition 
of oxygen or any other negative element, such as chlorine; 
reduction is the removal of oxygen or any other negative 
element. (&) From the standpoint of the ionic theory, 
oxidation is the addition of positive electricity, and reduc- 
tion the withdrawal. Thus, when ferrous chloride solution 
becomes ferric chloride upon the addition of chlorine, 
the ferrous ions (Fe ++ ) become ferric ions (Fe +++ ), while 
the electrically neutral chlorine molecules, having become 
ions, lose positive electricity, i.e. they undergo reduction, 
(c) Occasionally, oxidation is spoken of as occurring when 
the valence is raised, and reduction when the valence is 
lowered. For instance, in the change from ferrous to ferric 
compounds the valence of iron is raised from two to three, 
while by the reduction of ferric to ferrous the valence is low- 
ered from three to two. (Compare interaction of mercuric 
chloride and stannous chloride.) 

Oxides and Hydroxides of Iron. — Iron forms three oxides. 



546 INORGANIC CHEMISTRY 

Ferrous oxide (FeO) is an unstable black powder. Ferric 
oxide (Fe 2 3 ) occurs native in many varieties as hematite 
■ — the most abundant ore of iron. It can be prepared by 
heating ferrous sulphate or ferric hydroxide. Large quan- 
tities are manufactured from the ferrous sulphate obtained 
as a by-product in the cleansing of the iron used in making 
galvanized and tinned ware. It is sold under the names 
rouge, crocus, and Venetian red, and is used to polish glass 
and jewelry and to make red paint. Ferrous-ferric, or 
ferroso-ferric, oxide (magnetic oxide of iron, Fe 3 4 ), occurs 
native as magnetite ; if magnetic, it is called loadstone. 
It is produced as a film or scale by heating iron in the air. 
The firm coating of this oxide formed by exposing iron to 
steam protects the metal from further oxidation ; iron thus 
coated is called Russia iron. Some authorities call this oxide 
iron ferrite (Fe(Fe0 2 )2). Ferrous hydroxide (Fe(OH) 2 ) is a 
white solid formed by the interaction of a ferrous salt and an 
alkali, such as sodium hydroxide. Exposed to the air, it soon 
turns green, and finally brown, owing to the formation of 
ferric hydroxide. Ferric hydroxide (Fe(OH) 3 ) is a reddish 
brown solid, formed by the interaction of an hydroxide (e.g. 
sodium hydroxide) and a ferric salt. It readily forms a 
colloidal solution (compare Arsenic Trisulphide, page 404). 

Ferrous Sulphate, FeS0 4 , is a green salt obtained by the 
interaction of iron (or ferrous sulphide) and dilute sulphuric 
acid, and is a by-product in several large industries (e.g. 
see Ferric Oxide). It is also prepared on a large scale by 
oxidizing iron pyrites (FeS 2 ) . This is accomplished simply by 
roasting, or more often by exposing heaps of pyrites to moist 
air; the mass is extracted with water containing scrap iron 
and a small proportion of sulphuric acid, and large light 
green crystals are obtained from the solution. The crys- 
tallized salt (FeS0 4 . 7 H 2 0) is also called green vitriol or 



IRON 547 

copperas. Exposed to the air ; ferrous sulphate effloresces 
and oxidizes. Large quantities are used as a mordant in 
dyeing silk and wool, as a disinfectant and a wood preserva- 
tive, and in manufacturing ink ; bluing, pigments, leather, 
varnish, and mottled soaps. Much black writing ink is 
made essentially by mixing ferrous sulphate, nutgalls, gum, 
and water. A mixture of lime and ferrous sulphate is used 
as a coagulant in purifying water and sewage. 

Ferric Sulphate, Fe 2 (S0 4 ) 3 , is formed by oxidizing an 
acid solution of ferrous sulphate with nitric acid. When 
ferric sulphate solution is mixed with the proper quan- 
tity of potassium (or ammonium) sulphate, iron alum 
(K 2 S0 4 . Fe 2 (S0 4 ) 3 . 24 H 2 or K 2 Fe 2 (S0 4 ) 4 . 24 H 2 0) is formed. 
It is a pale violet, crystalline solid, which has properties 
like ordinary alum. Iron alum is used chiefly as a 
mordant. 

Iron Sulphides. — There are two iron sulphides. Com- 
mercial ferrous sulphide (FeS) is a black, brittle, metallic- 
looking solid, but the pure compound is yellow and crys- 
talline. It is also obtained as a fine black precipitate by 
the interaction of a dissolved ferric or ferrous salt and am- 
monium sulphide, though not by hydrogen sulphide. It is 
made on a large scale by fusing a mixture of iron and sulphur. 
Its chief use is in preparing hydrogen sulphide. Ferric 
sulphide (iron disulphide, iron pyrites, pyrite, FeS 2 ) is one 
of the commonest minerals. It is a lustrous, metallic, 
brass-yellow solid. Crystals of pyrites, found in many rocks, 
are often mistaken for gold — hence the popular name, 
" fool's gold." It is valueless as an iron ore, but large quan- 
tities are used as a source of sulphur in making sulphuric 
acid. Over one and a half million tons are annually con- 
sumed in the sulphuric acid industry. 



548 INORGANIC CHEMISTRY 

Iron Chlorides. — When iron interacts with hydrochloric 
acid, ferrous chloride (FeCl 2 ) is formed in solution. Ferric 
chloride (FeCl 3 ) is readily prepared by passing chlorine gas 
into a solution of ferrous chloride. It is a dark, lustrous, 
crystalline solid ; but owing to its extreme deliquescence, 
it is usually sold as a solution, which is a brown liquid. Nas- 
cent hydrogen or another reducing agent changes ferric chlo- 
ride into ferrous chloride. It hydrolyzes readily. 

Ferrous Carbonate, FeC0 3 , occurs native as the iron ore 
siderite, clay ironstone, or spathic iron ore. The typical 
variety is light yellow or brown, lustrous, crystalline, and 
not very hard; but many kinds are impure, and the prop- 
erties vary. It is slightly soluble in water containing 
carbon dioxide, and is therefore found in some mineral 
springs. (See Chalybeate Waters.) Like all carbonates, 
it yields carbon dioxide with warm hydrochloric acid. 

Iron Cyanides. — Iron and cyanogen (CN) 2 , with or with- 
out potassium, form several compounds. The most im- 
portant is potassium ferrocyanide (K 4 Fe(CN) 6 ). It is a 
lemon-yellow, crystalline solid, containing three molecules 
of water of crystallization. Unlike most cyanogen com- 
pounds, it is not poisonous. Its commercial name is yellow 
prussiate of potash. It is manufactured by fusing together 
iron filings, potassium carbonate, and nitrogenous animal 
matter (such as horn, hair, blood, feathers, and leather). 
The mass is extracted with water, and the salt is separated 
by crystallization. In Germany this salt is manufactured 
from the iron oxide which has been used to purify illuminat- 
ing gas. Large quantities are used in dyeing and calico 
printing, and in making bluing and potassium cyanogen 
compounds. Potassium ferricyanide (K 3 Fe(CN) 6 ) is a dark 
red, crystalline solid, containing no water of crystallization. 
It is often called red prussiate of potash. It is manufactured 



IRON 549 

by oxidizing potassium ferrocyanide with potassium per- 
manganate or chlorine, thus: — 

2 K 4 Fe(CN) 6 + CI, = 2 K 3 Fe(CN) 6 + 2 KC1 

Potassium Chlorine Potassium Potassium 

Ferrocyenide Ferricyanide Chloride 

It is very soluble in water, forming a yellow, unstable solu- 
tion. In alkaline solution it is a vigorous oxidizing agent, 
and finds extensive use in dyeing. It is also used as one of 
the ingredients of the sensitive coating of " blueprint" paper. 

The valence of the radical Fe(CN) 6 is three in ferri- 
cyanides and four in ferrocyanides. 

Ferrous salts and potassium ferricyanide interact in solu- 
tion and produce ferrous ferricyanide (Fe 3 (Fe(CN) 6 ) 2 ). 
This is a blue solid and is often called Turnbull's blue. But 
ferrous salts produce with potassium ferrocyanide a white 
precipitate (ferrous ferrocyanide) which quickly oxidizes 
to a complex blue compound. Ferric salts interact with 
potassium ferrocyanide and produce ferric ferrocyanide 
(Fe 4 (Fe(CN) 6 ) 3 ). This is a dark blue solid, and is called 
Prussian blue or Berlin blue. Ferric salts produce no pre- 
cipitate with potassium ferricyanide. Prussian blue is ex- 
tensively used in dyeing and calico printing, and in making 
bluing. The above reactions, which allow ferrous and 
ferric salts to be distinguished, may be summarized as 
follows : — - 





Cyanide 


Ferrous Salt 


Ferric Salt 


Ferrocyanide 
Ferricyanide 


Whitish precipitate 
Turnbull's blue 


Prussian blue 
No precipitate 





Besides the above tests, potassium sulphocyanate (KCNS) 
produces a blood-red solution of ferric sulphocyanate 
(Fe(CNS) 3 ) with ferric salts, but leaves ferrous salts un- 
changed. The tests for iron are thus numerous and specific. 



550 INORGANIC CHEMISTRY 

Miscellaneous. — Iron compounds yield several kinds of 
ions ; e.g. the ferrous ion (Fe ++ ) and the ferric ion (Fe ++ " r ) ; 
each is colorless, though the former appears delicate green 
and the latter very pale yellow, owing to traces of other sub- 
stances. Complex ions are Fe(CN) 6 and Fe(CN) 6 . 

The atomic weight of iron is 55.84. 

Nickel 

Nickel, Ni, occurs combined with arsenic or with arsenic 
and sulphur as niccolite (NiAs) and nickel glance (NiAsS). 
Small amounts of metallic nickel are found in meteorites. 
The ores which furnish most of the commercial nickel are 
the nickel-bearing iron sulphides of the Sudbury district, 
Canada, and the silicate of nickel and magnesium (gar- 
nierite) found in New Caledonia. 

Preparation and Properties. — Nickel is obtained from its 
ores by a complicated smelting or electrolytic process. It 
is a white, hard metal. It is ductile and tenacious. It 
takes a brilliant polish and does not tarnish in dry air, 
though in moist air it tarnishes very slowly. Like iron, it 
is attracted by a magnet. 

Uses of Nickel. — For many years it has been used as one 
ingredient of the small coins of several countries. The per 
cent of nickel varies from 12 in the United States cent to 
25 in the five-cent piece. German silver contains from 15 
to 25 per cent of nickel, the rest being copper and zinc. 
Large quantities of nickel are used to coat or plate other 
metals, especially iron and brass. The nickel plating is 
done by electrolysis, as in the case of silver and gold plating, 
though the electrolytic solution used is a sulphate of nickel 
and ammonium ((NH 4 ) 2 S0 4 . NiS0 4 ) — not a cyanide, as in 
the other cases. The deposit of nickel is hard, brilliant, 



NICKEL 551 

and durable. Nickel becomes malleable if a little magnesium 
or aluminium is added to the molten metal, and sheets of 
iron covered with such nickel are made into household 
utensils. Nickeloid is a nickel-plated sheet zinc. Its 
attractive appearance and non-corrosive property adapt it 
for the manufacture of reflectors, refrigerator linings, bath 
tubs, show cases, and signs. Nickel is used in the manu- 
facture of nickel steel. This contains varying proportions 
of nickel ; large quantities are used for burglar-proof safes, 
and the armor plates and turrets of battleships. An alloy 
of nickel and copper known as monel metal is used where a 
non-corroding metal is desirable. 

Compounds of Nickel. — Nickel forms two series of com- 
pounds — the nickelous and the nickelic. The valence of 
nickel is two in the nickelous compounds and three in the 
nickelic. The nickelous compounds are the more common. 
Many nickel salts are green. The apple-green nickelous 
hydroxide (Ni(OH) 2 ) is formed by the interaction of an 
alkali and a dissolved nickel salt. Nickelous sulphide (NiS) 
is obtained as a black precipitate by the interaction of 
ammonium sulphide and a dissolved nickel salt. Nickelous 
sulphate (NiS0 4 ) and nickelous chloride (NiCl 2 ) are the usual 
commercial salts. Nickel carbonyl (Ni(CO) 4 ) is prepared 
by passing carbon monoxide over finely divided nickel. 
It is a volatile, colorless liquid which boils at 43° C. The 
vapor is poisonous and decomposes at 150°-180° C. into 
metallic nickel and carbon monoxide. This compound has 
a technical application in the Monde process of preparing 
nickel. Ammonium nickelous sulphate ((HN^SCX • NiS0 4 ) 
was mentioned in the preceding paragraph. 

Tests for Nickel. — The formation of the green nickelous 
hydroxide is a distinctive test. Nickel compounds color 



552 INORGANIC CHEMISTRY 

a borax bead brown in the oxidizing flame and gray in the 
reducing flame. 

Miscellaneous. — Nickelous salts yield a green ion (Ni ++ ). 
The atomic weight of nickel is 58.68. 

Cobalt 

Cobalt, Co, generally occurs combined with arsenic or 
arsenic and sulphur as smaltite (CoAs 2 ) and cobaltite 
(CoAsS), which are usually associated with the correspond- 
ing nickel compounds. 

Preparation and Properties. — Cobalt is obtained as a 
powder by reducing its oxide with hydrogen or as a coherent 
mass by the aluminothermic method. It is a white metal 
with a faint reddish tinge. Like nickel, it takes a brilliant 
polish. It is less magnetic than iron. Metallic cobalt has 
few uses. 

Compounds of Cobalt. — Cobalt, like nickel, forms two 
series of compounds — the cobaltous and the cobaltic. The 
valence of cobalt is two in cobaltous compounds and three 
in cobaltic. The cobaltous compounds are the more com- 
mon, though many complex cobalt compounds are known. 
Cobaltous sulphide (CoS) is obtained as a black precipi- 
tate by the interaction of ammonium sulphide and. a dis- 
solved cobalt salt. Cobaltous chloride (CoCl 2 ) and cobaltous 
nitrate (Co(N0 3 ) 2 ) are red crystalline salts; they crystallize 
with six molecules of water of crystallization. When part 
or all of the water of crystallization is driven off, these salts 
become violet or blue. A complex cobalt compound known 
as smalt, or smalt blue, is used to decorate porcelain. It 
has a variable composition, but is essentially a cobalt silicate. 
Cobalt compounds color glass blue. Such glass transmits 
red, blue, and violet light, but not yellow, orange, and green, 



i 



COBALT 553 

so it is sometimes used to detect the violet potassium flame 
when masked by the yellow sodium flame. 

Tests for Cobalt. — Cobalt compounds color a borax bead 
blue in both flames. Cobaltous compounds when mixed 
with potassium nitrite and acetic acid form a yellow-white 
precipitate of potassium cobaltini trite (K 3 Co(N0 2 )6), thereby 
distinguishing cobalt from nickel. 

Problems and Exercises 

1. What is the simplest formula of a compound 9 gm. of which 
yielded 4.8 gm. of sulphur and 4.2 gm. of iron? 

2. (a) An iron oxide contains 27.6 per cent of oxygen and has 
the molecular weight of 232. What is its simplest formula? (b) If 
1.013 gm. of iron form 1.446 gm. of an iron oxide, what is the sim- 
plest formula of the oxide ? 

3. Calculate the approximate atomic weight of iron from the 
specific heat of the metal. 

4. What weight of oxygen is required to unite with 21 gm. of iron 
to give the magnetic oxide of iron ? What volume will this oxygen 
occupy at 10° C. and 750 mm. ? 

5. What weight of crystallized ammonium iron alum will be 
formed by the interaction of solutions containing 12 gm. of ammo- 
nium sulphate and 30 gm. of ferric sulphate ? 

6. How much coke (containing 96.5 per cent of carbon) is needed 
to reduce 175 tons of hematite (96.5 per cent pure) ? 

7. Calculate the weight of cobalt or nickel in (a) 27 gm. of 
potassium cobaltini trite, K 3 Co(N0 2 )6; (&) 52 gm. of cobaltous 
nitrate; (c) 100 milligrams of nickel carbonyl, Ni(CO) 4 ; (d) 235 
centigrams of ammonium nickelous sulphate, (NH 4 ) 2 S0 4 . NiS0 4 . 

8. Write the formulas of the following compounds and indicate 
the valence of each element : Potassium ferrocyanide, potassium 
ferricyanide, ferric ferrocyanide, ferrous ferricyanide, potassium 
cobaltinitrite. 

9. How many cc. of ammonia solution having a specific gravity of 
.9605 and containing 9.83 per cent of NH 3 by weight are needed to 
precipitate the iron as Fe(OH) 3 from 1 gm. of (NH 4 ) 2 S0 4 . FeS0 4 . 
6H 2 0? 



CHAPTER XXXIV 
Platinum and Associated Metals 

Occurrence of Platinum. — Platinum occurs as the essen- 
tial ingredient of platinum ore or so-called native platinum. 
The ore contains from 60 to 86 per cent of platinum. The 
other metals present are ruthenium, osmium, iridium, 
rhodium, and palladium. Iron, gold, and copper are also 
usually present. Only one native compound is known, 
viz. platinum arsenide (sperrylite, PtAs 2 ). These metals, 
like gold, are noble metals, i.e. they do not unite with oxygen 
at any temperature. 

The word platinum is derived from platina, a form of the Spanish 
word plata, meaning silver, because native platinum was regarded 
as an impure ore of silver by the Spaniards, who first discovered it 
in South America about 1735. Platinum is now sometimes called 
by its old name -platina. 

Preparation of Platinum. — The platinum ore, which 
occurs as rounded grains .or flattened scales in alluvial de- 
posits, is first digested with mercury or dilute aqua regia 
to remove the gold, silver, and copper; and then with con- 
centrated aqua regia, which changes all the platinum and 
a very little iridium into soluble compounds, leaving behind 
an alloy of iridium and osmium. From the clear solution 
the platinum and iridium are precipitated by ammonium 
chloride as compounds, which, on heating, yield the metals 
as a spongy mass. This spongy platinum is melted in a 
lime crucible with an oxyhydrogen flame or in the electric 
furnace, and hammered while hot into sheet platinum. 

554 



TLATINUM AND ASSOCIATED METALS 555 

The very small amount of iridium is seldom removed from 
the metallic platinum. 

Properties and Uses of Platinum. — Platinum is a lustrous, 
gray-white, soft metal. It is malleable and ductile, and 
usually appears in commerce in the form of wire, sheet, and 
dishes. Sheet platinum is cut into squares — the familiar 
platinum foil of the laboratory, or made into crucibles, 
dishes, and stills. Its use in these forms is due partly to 
its infusibility and partly to its resistance to acids and other 
corrosive chemicals. Although it is attacked by fused 
caustic alkalies, low melting metals, aqua regia, and a few 
other substances, it is practically indispensable in the chemi- 
cal laboratory and is used in many chemical processes in- 
volving accurate analysis. Platinum is a good conductor of 
electricity, and large quantities are consumed in electrical 
apparatus, especially incandescent electric light bulbs. 
Short pieces of wire are fused through the glass at the base of 
the bulb and thereby serve as electric conductors. Platinum 
is the only metal thus far found which is perfectly adapted 
to this use, because it has the same coefficient of expansion 
as glass. Recently platinum has come into use as jewelry. 
Platinum has a specific gravity of about 21, which is higher 
than that of any known substance, except osmium and irid- 
ium. The melting point is about 1755° C. In the form of 
a black, porous mass it is called spongy platinum, and a still 
finer form is called platinum black. Both forms absorb 
large volumes of gases ; and if a current of gas is directed 
against the metal, the gas often takes fire. Coherent plati- 
num has the same property to a less degree, for it becomes 
red-hot if held in a stream of illuminating gas, and often ig- 
nites the gas. Finely divided platinum is used as the cata- 
lyzer in the contact method for making sulphuric acid. Plati- 
num forms alloys with other metals, and should never be 



556 INORGANIC CHEMISTRY 

heated with lead, similar metals, or their compounds, since 
the alloys have a low melting point. With iridium, however, 
it forms a very hard alloy of which the standard meters are 
made. 

Compounds of Platinum. — Chloroplatinic acid (H 2 PtCl 6 ) 
is a brownish, deliquescent solid prepared by treating 
platinum with aqua regia and evaporating the solution to 
crystallization. By carefully heating chloroplatinic acid 
in a current of chlorine, platinum tetrachloride (PtCl 4 ) is 
obtained ; it resembles chloroplatinic acid in color but is 
not deliquescent. The acid forms salts, the best known 
being the yellow, crystalline potassium chloroplatinate 
(K 2 PtCl 6 ) and ammonium chloroplatinate ((NH 4 )2PtCl 6 ). 

Palladium is used in chemical analysis to absorb hydrogen, in 
making scientific instruments, as a catalyst, and as a substitute for 
platinum. A native (as well as an artificial) alloy of iridium and os- 
mium, called iridosmine, is used to tip gold pens. Iridosmine is 
often called osmiridium. 



Problems and Exercises (Review) 

1. (a) How much sodium chloride is needed to prepare a kilo- 
gram of hydrogen chloride ? (b) A kiloliter at 20° C. and 765 mm. ? 
(c) What weight of chlorine can be obtained from the gas prepared 
in (6) ? 

2. Find the weight of a mixture of 210 cc. of oxygen and 790 cc. 
of nitrogen. (Standard conditions.) 

3. Discuss the following topics : (a) Electrolytes depress the 
freezing point abnormally ; (b) ions migrate to their respective elec- 
trodes ; (c) tests are for ions. 

4. (a) The equivalent weight of zinc was found by experiment to 
be 32.6 and the specific heat .095. What is the approximate atomic 
weight? (b) Similarly for tin, equiv. wt. = 29.42 and sp. ht. = 
.056. Calculate the approximate atomic weight. 



APPENDIX 

i. The Metric System of weights and measures is used in physics 
and chemistry. 

It is based on the meter. This is the unit of length, and it is a little 
longer than a yard. Its exact length is 39.37 inches — a number to 
remember. Lengths shorter than a meter are called decimeters, 
centimeters, and millimeters. Deci- means .1, centi- means .01, and 
milli- means .001. Lengths longer than a meter are called deca- 
meters, hectometers, and kilometers. Deca- means 10, hecto- means 
100, kilo- means 1000. Notice that all these relations are decimal. 
Hence, a meter contains 10 decimeters, 100 centimeters, or 1000 
millimeters. It is also evident that 10 millimeters equal 1 centi- 
meter, 10 centimeters equal 1 decimeter, and that 1000 meters equal 
1 kilometer. The millimeter, centimeter, decimeter, and meter are 
the denominations most frequently used in physical science to express 
length, though very long distances are expressed in kilometers. It 
is advisable to remember that — 

1 decimeter = about 4 inches. 
30 centimeters = about 1 foot. 
2.5 centimeters = about 1 inch. 

The customary abbreviations of the linear denominations are 
meter, m.; decimeter, dm.; centimeter, cm.; and millimeter, mm. 

The unit of weight is the gram. It is a small weight, being only 
about one thirtieth of an ounce. A five-cent coin weighs approxi- 
mately five grams. The weights of small objects and the small 
quantities used in chemical analysis are expressed in terms of the 
gram. The weights of heavy objects and large quantities are often 
expressed in terms of the kilogram, which is 1000 times heavier than 
the gram. Just as the meter is subdivided, so the gram is subdivided 
decimally into smaller weights called the decigram, centigram, and 
milligram. Therefore, a gram contains 10 decigrams, 100 centi- 
grams, or 1000 milligrams; and a kilogram contains 1000 grams. 

557 



558 INORGANIC CHEMISTRY 

The gram, decigram, centigram, milligram, and occasionally the 
kilogram, are used in physical science, though the gram is the most 
common denomination. For example, if an object weighs 2 grams, 
2 centigrams, and 5 milligrams, the weight is expressed as 2.025 
grams; similarly, 5 milligrams is often expressed as .005 gram. 
Preferable abbreviations of the weight denominations are gram, gm.; 
decigram, dg.; centigram, eg.; milligram, mg.; and kilogram, kg. 

The unit of volume is the liter. It is slightly larger than a quart, 
and is used for both dry and liquid measure. As in the case of the 
meter and the gram, the liter is subdivided into denominations called 
the deciliter, etc. But these fractional denominations are seldom 
used. That is, small volumes are not expressed as decimal fractions 
of a liter, but as cubic centimeters. A liter contains 1000 cubic 
centimeters, and parts of a liter are designated by the proper num- 
ber of cubic centimeters. For example, one half a liter is called 500 
cubic centimeters, one fourth is 250 cubic centimeters, one tenth is 
100 cubic centimeters; two liters is often called 2000 cubic centi- 
meters. The relation of cubic centimeters to a liter is simple. The 
French chemists who devised the metric system first found the length 
of the meter by measuring a part of the meridian passing near Paris. 
Subsequently, they constructed a vessel equal to the capacity of a 
cubical vessel having edges 10 centimeters long; such a cubical vessel 
contains, of course, 1000 cubic centimeters. The capacity of this 
vessel they named the liter. Therefore, the liter and 1000 cubic 
centimeters are identical, whatever the substance measured — a fact 
to remember. The abbreviation of liter is 1., and of cubic centimeter 
is cc. or cm. 3 . 

The relation between meter and liter has been shown. Another 
important relation should be noted. The cubical vessel named the 
liter may, of course, be filled with any substance; if it is filled with 
pure water at 4° C, that weight of water is called a kilogram. There- 
fore in the case of water the following relation exists: 1 liter, 1 quart, 
and 1000 cubic centimeters weigh approximately the same as 1 kilo- 
gram, 1000 grams, and 2.2 pounds. Since many liquids have about 
the same specific gravity as water, this general relation is useful, and 
should be learned. It is clear from the relation just given that 1 
cubic centimeter of water weighs 1 gram — a fact to remember, since 
this relation enables us to convert volume into weight, and vice versa. 

The relation between the units, multiples, and submultiples of the 
metric system is shown in the — 



APPENDIX 



559 



Table of the Metric System 



Length 


Weight 


Volume 


Notation 


Kilometer 


Kilogram 


Kiloliter 


1000. 


Hectometer 


Hectogram 


Hectoliter 


100. 


Decameter 


Decagram 


Decaliter 


10. 


Meter 


Gram 


Liter 


1. 


Decimeter 


Decigram 


Deciliter 


0.1 


Centimeter 


Centigram 


Centiliter 


0.01 


Millimeter 


Milligram 


Milliliter 


0.001 



From this table it is evident that 10 milligrams equal 1 centigram, 
10 centigrams equal 1 decigram, 10 decigrams equal 1 gram, and 
so on. 

The relation of the metric system to weights and measures in com- 
mon use is shown by the — 

Table of Metric Equivalents 



1 meter 

1 kilometer 
1 centimeter 

1 liter 
1 liter 



39.37 inches 

0.62 mile 
0.39 inch 

= 0.908 quart (dry) 
= 1.056 quarts (liq.) 



1 gram = 15.452 grains 

1 kilogram =2.2 pounds 

(avoir.) 
1 metric ton = 2204 pounds 



1 inch 

1 mile 

1 cubic inch 

1 quart (liq.) 
1 pound (avoir.) 

1 ounce (avoir.) 
1 ounce (troy) 



2.54 centime- 
ters 
1.6 kilometers 
16.39 cubic 
centimeters 
0.9465 liter 
0.4536 kilo- 
gram 
28.35 grams 
31.1 grams 



1 grain (apoth.) = 0.0648 gram 



The passage from the English to the metric system may be accom- 
plished by utilizing the — 

Table of Metric Transformation 



To Change 

t 

Inches to centimeters .... 
Centimeters to inches .... 
Cubic inches to cubic centimeters 
Cubic centimeters to cubic inches 
Ounces to grams (avoir.) . . . 
Grams to ounces (avoir.) . . . 

Grains to grams 

Grams to grains 



Multiply by 



2.54 

0.3937 
16.387 

0.061 
28.35 

0.0353 

0.0648 
15.43 






560 



INORGANIC CHEMISTRY 



Problems 

1. What is the abbreviation of gram, centigram, liter, meter, cubic 
centimeter, centimeter, decimeter, milligram? 

2. Express (a) 1 liter in cubic centimeters, (b) 2 1. in cc, (c) 1 meter 
in centimeters, (d) 250 cm. in dm., (e) 1 kg. in grams, (/) 250 gm. 
in mg. 

3. Add 2 kg., 5 dg., 2 eg., 4 gm., and 7 mg., and express the sum 
in grams. 

4. How many cc. in a liter? 

5. What is the weight in grams of (a) 1 liter of water, (b) 250 cc, 
(c) 500 cc, (d) 721 cc? 

6. Express in grams (a) 721 kg., (b) 62 mg., (c) 245 eg., (d) 84 dg. 

7. Express (a) 40 meters in inches, (6) 25 kilograms in pounds, 
(c) 54 grams in ounces, (d) 72 grams in grains, (e) 75 liters in quarts 
(Hq.). 



too 



212 



2. The Thermometer in scientific use is the centigrade. The boil- 
ing point of water on this thermometer is 100, and the freezing point 
is (Fig. 83). The equal spaces between these points are called 
degrees. The abbreviation for centigrade is C, and for degrees is °. 
Thus, the boiling point of water is 100° C. Degrees below zero are al- 
ways designated as minus, e.g. — 12° C. means 12 degrees below zero. 
The thermometer in popular use is the Fahrenheit. 
On this instrument the boiling point of water is 
212° and the freezing point is 32° above zero (Fig. 
83). The abbreviation for Fahrenheit is F. 

To change Fahrenheit degrees into the equiva- 
lent centigrade degrees, subtract 32 and multiply 
the remainder by f , or briefly — 

C. = f (F.-32). 

To change centigrade degrees into the equivalent 
Fahrenheit temperature, multiply by f and add 32 
to the product, or briefly — 

F. = fC.4-32. 

The point -273°C. is called absolute zero. Abso- 
lute temperature is reckoned from this point. 
Degrees on the absolute scale are found by adding 

273 to the readings on the centigrade thermometer. Thus, 273° 

absolute is 0° C, 274° absolute is + 1° C, etc. 




Fig. 83. — Ther- 
mometers. 



APPENDIX 



561 



Problems 

1. Change into Fahrenheit readings the following centigrade read- 
ings: (a) 60.5, (b) 40, (c) 92, (d) -5, (e) 0, (f) 100, (g) 860, (h) -40. 

2. Change into centigrade readings the following Fahrenheit read- 
ings: (a) 207, (6) 180, (c) 0, (d) -30, (e) 212, (/) 100, (g) -40, (ft) 270. 

3. Express the following centigrade readings in absolute readings: 
(a) 0, (b) 24, (c) -13, (d) -260. 

3. Crystallization. — Most substances in passing from a liquid or 
a gas into a solid assume a definite shape. This change is called 
crystallization, and the substances are said to crystallize or to form 
crystals. Crystals are produced by (1) evaporating a solution, 
(2) cooling a melted solid, or (3) cooling a vapor. Thus, sodium 
chloride crystals are formed by evaporating a salt solution; sulphur 
crystals, by melting and then cooling sulphur; and iodine crystals, 
by heating iodine in a test tube. These methods are called, respec- 
tively, evaporation, fusion, and sublimation. 

As a rule, each substance has an individual crystal form by which 
it can be distinguished. Although there are thousands of different 
crystals, all belong to one of six classes or systems. This classifica- 
tion is based upon two assumptions: (1) all crystals contain certain 
lines called axes, and (2) the surfaces or faces are grouped around the 
axes in definite positions. The axes connect angles, edges, or faces, 
which are similarly situated on opposite sides of the crystal. The 
bounding planes or faces are arranged symmetrically around the axes, 
which also determine (by their lengths and relative positions) the 
positions of the bounding planes. For example, the cube has three 
equal axes at right angles to one another and terminating in the 
center of each of the six bounding surfaces. 

The following is a brief description of the six systems of crystal- 
lization: — 





Fig. 84. — Isometric crystals (cube, octahedron, dodecahedron). 



(1) Isometric. — This has three equal axes intersecting at right 
angles. The simplest forms are the cube, octahedron, and dodeca- 



562 



INORGANIC CHEMISTRY 



hedron (Fig. 84). Substances crystallizing in this system are dia- 
mond, common salt, alum, fiuor spar, iron pyrites, and garnet. 



\KJ 




Fig. 85. — Tetragonal crystals. 



(2) Tetragonal. — This has three axes at right angles; but one 
axis is shorter or longer than the other two, which are equal. The 
common forms are the prism, pyramid, and their combinations 
(Fig. 85). Tin dioxide and zircon form tetragonal crystals. 




Fig. 86. — Orthorhombic crystals. 

(3) Orthorhombic. — This has three unequal axes intersecting at 
right angles. Common forms are the prism, pyramid, and their com- 
binations (Fig. 86). Potassium nitrate, barium sulphate, topaz, and 
native sulphur crystallize in this system. 



<1 




J>l 


^,i 




i N 



A 




/ 




\ v 


| 









.J 


7 




Fig. 87, — Hexagonal crystals. 

(4) Hexagonal. — This has four axes: three are equal and inter- 
sect at 60° in the same plane; the fourth is longer or shorter than 
the others and is at right angles to their plajxe, \\ is a complex 



APPENDIX 



563 



system. Common forms are the prism, pyramid, rhombohedron, 
scalenohedron, and their combinations (Fig. 87). In this extensive 
system are found quartz, calcite, beryl, corundum, and ice (see Figs. 
60 and 71). 

(5) Monoclinic. — This has three unequal axes: two cut each other 
obliquely, and the third is at right angles to the plane of the other 
two. Common forms are combinations of prisms. It is a complex 
system, but includes many substances, e.g. sulphur deposited by 
fusion, sodium carbonate, borax, gypsum, and ferrous sulphate 
(Fig. 88). 






Fig. 88. — Monoclinic crystal. 



Fig. 89. — Triclinic crystals. 



(6) Triclinic. — This has three unequal axes, all intersecting at 
oblique angles. Common forms are complex combinations. Copper 
sulphate, potassium dichromate, boric acid, and several minerals 
form triclinic crystals (Fig. 89). 



4. Vapor Pressure. • 
can be found in the 



The value of a in the formula in Chapter V 



Table of Vapor Pressure 



t 


a 


t 


(/ 


t 


a 


t 


a 


10 


9.18 


16 


13.57 


22 


19.66 


28 


28.10 


.5 


9.49 


.5 


14.00 


.5 


20.26 


.5 


28.93 


11 


9.81 


17 


14.45 


23 


20.88 


29 


29.79 


.5 


10.14 


.5 


14.91 


.5 


21.52 


.5 


30.66 


12 


10.48 


18 


15.38 


24 


22.18 


30 


31.56 


.5 


10.83 


.5 


15.87 


.5 


22.85 


.5 


32.47 


13 


11.19 


19 


16.37 


25 


23.55 


31 


33.41 


.5 


11.56 


.5 


16.88 


.5 


24.26 


.5 


34.37 


14 


11.94 


20 


17.41 


26 


24.99 


32 


35.36 


.5 


12.33 


.5 


17.95 


.5 


25.74 


.5 


36.37 


15 


12.73 


21 


18.50 


27 


26.51 


33 


37.41 


.5 


13.14 


.5 


19.07 


.5 


27.29 


.5 


38.47 



564 



INORGANIC CHEMISTRY 



5. Atomic Weights. — The following is a table of 

International Atomic Weights (1916) 



Element 


Sym- 
bol 


At. Wt. 


Element 


Sym- 
bol 


At. Wt. 


Aluminium . . 


Al 


27.1 


Molybdenum . 


Mo 


96.0 


Antimony . 






Sb 


120.2 


Neodymium . 




Nd 


144.3 


Argon . . 






A 


39.88 


Neon . . . 




Ne 


20.2 


Arsenic . °. 






As 


74.96 


Nickel . 






Ni 


58.68 


Barium . 






Ba 


137.37 


Niton . 






Nt 


222.4 


Bismuth 






Bi 


208.0 


Nitrogen 






N 


14.01 


Boron . . 






B 


11.0 


Osmium 






Os 


190.9 


Bromine 






Br 


79.92 


Oxygen 






O 


16.00 


Cadmium . 






Cd 


112.40 


Palladium 






Pd 


106.7 


Caesium . . 






Cs 


132.81 


Phosphorus 






P 


31.04 


Calcium 






Ca 


40.07 


Platinum 






Pt 


195.2 


Carbon . . 






C 


12.005 


Potassium 






K 


39.10 


Cerium . . 






Ce 


140.25 


Praseodymium 




Pr 


140.9 


Chlorine 






CI 


35.46 


Radium 






Ra 


226.0 


Chromium . 






Cr 


52.0 


Rhodium 






Rh 


102.9 


Cobalt . . 






Co 


58.97 


Rubidium 






Rb 


85.45 


Columbium x 






Cb 


93.5 


Ruthenium 




Ru 


101.7 


Copper . . 






Cu 


63.57 


Samarium 




Sa 


150.4 


Dysprosium 






Dy 


162.5 


Scandium . . 




Sc 


44.1 


Erbium . . 






Er 


167.7 


Selenium 






Se 


79.2 


Europium . 






Eu 


152.0 


Silicon . 






Si 


28.3 


Eluorine 






F 


19.0 


Silver . 






Ag 


107.88 


Gadolinium 






Gd 


157.3 


Sodium 






Na 


23.00 


Gallium . . 






Ga 


69.9 


Strontium 






Sr 


87.63 


Germanium 






Ge 


72.5 


Sulphur 






S 


32.06 


Glucinum 2 






Gl 


9.1 


Tantalum 






Ta 


181.5 


Gold . . 






Au 


197.2 


Tellurium 






Te 


127.5 


Helium . . 






He 


4.00 


Terbium 






Tb 


159.2 


Holmium . 






Ho 


163.5 


Thallium 






Tl 


204.0 


Hydrogen . 






H 


1.008 


Thorium 






Th 


232.4 


Indium . . 






In 


114.8 


Thulium 






Tm 


168.5 


Iodine . . 






I 


126.92 


Tin . . . 






Sn 


118.7 


Iridium . . 






Ir 


193.1 


Titauium 






Ti 


48.1 


Iron . . . 






Fe 


55.84 


Tungsten 






W 


184.0 


Krypton 






Kr 


82.92 


Uranium 






U 


238.2 


Lanthanum 






La 


139.0 


Vanadium 






V 


51.0 


Lead . . . 






Pb 


207.20 


Xenon . 






Xe 


130.2 


Lithium . . 






Li 


6.94 


Ytterbium 






Lutecium . 






Lu 


175.0 


(Neoytterbium) 


Yb 


173.5 


Magnesium 






Mg 


24.32 


Yttrium . . . 


Yt 


88.7 


Manganese . 






Mn 


54.93 


Zinc 


Zn 


65.37 


Mercury 






Hg 


200.6 


Zirconium .. . . 


Zr 


90.6 



1 Or Niobium, Nb. 



2 Or Beryllium, Be. 



INDEX 



Absolute alcohol, 314. 

And centigrade, 43, 44. 

Temperature, 43. 

Zero, 44, 560. 
Acetates, 319, 320, 450, 515. 
Acetic acid, 317. 

Glacial, 318. 

Ions, 318. 

Test, 320. 
Acetone, 316. 
Acetylene, 290. 

And calcium carbide, 285. 

As illuminant, 291. 

Burner, 292. 

Composition, 291. 

Explosion, 290. 

Flame, 291, 292. 

Generation, 292. 

Thermal properties, 292. 
Acid, Acetic, 317. 

Boracic, 374. 

Boric, 374. 

Butyric, 319, 322. 

Carbolic, 329. 

Carbonic, 280. 

Chamber, 348. 

Chloroplatinic, 556. 

Citric, 320. 

Disulphuric, 352. 

Fluosilicic, 384. 

Fuming nitric, 223. 

Fuming sulphuric, 351. 

Glacial acetic, 318. 

Glacial phosphoric, 395. 

Hydriodic, 189, 190, 371. 

Hydro bromic, 368. 

Hydrochloric, 203, 421. 

Hydrocyanic, 330. 

Hydrofluoric, 364. 

Hydro fluosilicic, 384. 

Hydro sulphuric, 337. 



Acid, Hypochlorous, 200, 201, 208. 

Lactic, 319, 326, 423. 

Manganic, 520. 

Metaboric, 375. 

Metaphosphoric, 395. 

Metasilicic, 381, 383. 

Muriatic, 203. 

Nitric, 216. 

Nitrous, 221. 

Oleic, 319. 

Orthophosphoric, 395. 

Orthosilicic, 381, 383. 

Oxalic, 319. 

Palmitic, 319, 321, 322, 323. 

Prussic, 330. 

Pyroligneous, 317. 

Pyrophosphoric, 396. 

Pyrosulphuric, 352. 

Silicic, 380, 381, 382. 

Stearic, 319, 321, 322, 323. 

Sulphuric, 345. 

Sulphurous, 342, 343. 

Tartaric, 319. 

Tetraboric, 375. 
Acid anhydrides, 163. 

Calcium carbonate, 281. 

Denned, 150. 

Dibasic, 157. 

Monobasic, 157. 

Of air, 282. 

Oxides, 162. 

Phosphate, 396. 

Potassium fluoride, 365. 

Potassium tartrate, 319, 320. 

Reaction, 150. 

Salt, 157, 165. 

Sodium carbonate, 166, 422. 

Sodium sulphate, 166. 

Sulphates, 351. 

Tribasic, 158. 
Acidity, 158. 
565 



566 



INDEX 



Acids, 149. 

Acetic series, 317. 

And hydrogen ions, 150. 

And non-metals, 150. 

And oxides, 162. 

Chlorine, 160. 

Composition, 150. 

Dissociation, degree, 164. 

Fatty series, 317. 

Formulas, 259. 

Hydrogen ions, 146. 

In butter, 322. 

Nomenclature, 160. 

Organic, 317, 319. 

Phosphorus, 395. 
Actinium, 531. 
Adsorption, 271. 
Air, 115. 

Acid of, 282. 

Alkaline, 213. 

And atmosphere, 115. 

And carbon dioxide, 121. 

And water vapor, 120. 

Bad, 120. 

Composition, 118. 

Fixed, 282. 

Liquid, 17, 125. 

Mixture, 124. 

Solubility, 65, 125. 

Weight of liter, 115. 
Albumin, 397. 
Alchemists, 458. 
Alcohol, 314. 

Absolute, 314. 

Denatured, 314. 

Ethyl, 312, 314. 

Methyl, 314. 

Test, 320. 

Triacid, 321. 

Wood, 314. 
Alcoholic fermentation, 315. 

Liquors, 315. 
Alcohols, 312. 
Aldehyde, formic, 316. 
Alizarin, 330. 
Alkali, 151, 422. 

Family, 435. 

Fixed, 161. 

Metals, 417, 435. 

Soda, 422. 
Alkalies, 151, 161. 



Alkalies. See Alkali and Base. 
Alkaline air, 213. 

Earth family, 475. 

Reaction, 151. 
Allotrope, 274. 

Allotropism, 274, 336, 387, 393. 
Alloys, 416. 

Aluminium, 447, 496. 

Antimony, 512. 

Copper, 446. 

Fusible, 407. 

Lead, 512. 

Manganese, 517. 

Platinum, 555. 

Silver, 455. 

Tin, 508. 

Zinc, 484. 
Alpha particles, 529, 531. 
Alum, 500. 

Cake, 499. 

Chrome, 525. 

Composition, 501. 

History, 501. 

Ions, 501. 

Iron, 547. 
Alumina, 496. 

Aluminates, 494, 496, 497, 498. 
Aluminium, 493. 

Acetate, 319, 501. 

Alloys, 447, 496. 

And radium, 529, 530. 

Bronze, 447. 

Carbide, 289. 

Chloride, 502. 

Compounds, hydrolysis, 499, 
501, 502. 

Electrolytic manufacture, 493, 
494. 

Gems, 497. 

Hydroxide, 497. 

In crust, 11, 493. 

Ions, 503. 

Metal and non-metal, 502. 

Occurrence, 493. 

Oxide, 496. 

Phosphate, 497. 

Properties, 494. 

Reduction by, 416, 495. 

Silicate, 503. 

Sulphate, 499. 

Test, 502. 



INDEX 



567 



Alumino-thermic method, 496. 
Aluminum. See Aluminium. 
Alundum, 496, 488. 
Amalgam, 416, 488. 

Gold, 488. 

Sodium, 419. 

Tin, 508. 

Zinc, 488. 
Amalgamation, 414, 453. 
Amethyst, 379, 497. 
Ammonia, 210, 213. 

Anhydrous, 212. 

Composition, 216. 

Copper compounds, 450. 

Dissociation, 212. 

Formation, 210, 219. 

Liquid, 212, 215. 

Muriate, 436. 

Of commerce, 211. 

Preparation, 210. 

Properties, 211. 

Refrigerant, 215. 

See Ammonium hydroxide. 

Soda process, 422. 

Water, 213. 
Ammoniacal liquor, 211, 297, 299. 
Ammonium, 214, 438. 

Carbonate, 438. 

Chloride, 435, 436. 

Chloroplatinate, 556. 

Compounds, 213, 214, 435. 

Hydroxide, 211, 213, 214, 215. 

Ion, 438. 

Molybdate, 527. 

Nickelous sulphate, 551. 

Nitrate, 221, 437. 

Polysulphide, 437. 

Sulphate, 437. 

Sulphide, 437. 

Sulphide, yellow, 437. 
Ammo no-compounds, 491. 
Amorphous, 266. 

Carbon, 266. 

Sulphur, 336. 
Amyl acetate, 321. 

Valerate, 321. 
Analysis, qualitative, 78, 340. 

Quantitative, 78. 

Spectrum, 438. 

Water, 54. 
An atmosphere, 116. 



Anglesite, 510. 

Anhydride, 163, 224, 280, 413. 
Anhydrite, 470. 
Anhydrous, 72. 

Copper sulphate, 72. 
Aniline, 329. 
Animal charcoal, 325. 
Anion, defined, 134. 
Anode, 141, 180. 
Anthracene, 330. 
Anthracite coal, 268. 
Antichlor, 203, 344. 
Antimony, 404. 

Alloys, 512. 

Compounds, 405, 406. 

Metalloid, 410. 

Name, 405. 

Oxides, 405 

Oxy chloride, 406. 

Test, 406. 
Apatite, 362, 391. 
Aqua ammonia, 211, 213. 

Fortis, 219. 

Regia, 224, 460. 
Argol, 319. 
Argon, 122, 441. 

In air, 118. 

Monatomic, 251. 
Aristotle, 115. 
Arrhenius, 133. 
Arsenic, 401. 

Acids, 403. 

Antidote, 402, 546. 

Marsh's test, 404. 

Metalloid, 410. 

Oxide, 402. 

Poisoning, 402. 

Properties, 402. 

Pyrites, 401. 

Salts, 403. 

Sulphides, 403. 

Sulpho-salts, 404. 

Test, 404. 

Trichloride, 402. 

Trioxide, 402. 

White, 402. 
Arsenious oxide, 402. 
Arsine, 404. 
Artificial diamonds, 264. 

Graphite, 266. 

Stone, 383. 



568 



INDEX 



Asbestos, 479. 
Ash, coal, 268, 270. 

Potassium compounds, 434. 

Silica, 380. 
Assaying, 453, 454. 
Assimilation of nitrogen, 131. 
Atmosphere, 115. 

An, 116. 

Ingredients, 116. 

Pressure, 116. 

See Air. 
Atomic groups, ammonium, 214. 

Combinations, 253, 257. 

Displacement, 257. 

Ions, 134. 

Valence, 255. 
Atomic theory, 91, 227, 233, 234. 
Atomic weights, 96. 

And equivalent weights, 231, 
244. 

And properties, 357. 

And specific heat, 246. 

And symbols, 97. 

And valence, 254. 

Approximate, 245. 

Carbon, 243. 

Determination, 228, 242-248. 

Exact, 246, 247. 

International table, 564. 

Oxygen, 243. 

Standard, 228, 237. 

Zinc, 247, 248. 
Atoms, 92. 

And ions, 134. 

And molecules, 92, 95. 

Combination, 253, 257. 

Decay, 530, 531. 

Decomposition, 95, 530, 531. 

Displacement, 258. 
Atoms in molecule, 250. 

Argon, 251. 

Arsenic, 402. 

Arsenious oxide, 403. 

Bromine, 367. 

Cadmium, 251, 486. 

Chlorine, 251. 

Fluorine, 363. 

Hydrofluoric acid, 365. 

Hydrogen, 236, 251. 

Iodine, 370. 

Mercury, 251, 488. 



Atoms, Nitrogen, 251. 

Oxygen, 237, 251. 

Ozone, 251. 

Phosphorus, 251, 394. 

Potassium, 251. 

Sodium, 251, 419. 

Sulphur, 251, 335. 

Zinc, 251, 483. 
Attraction and repulsion, 141. 
Auric chloride, 461. 
Avogadro's hypothesis, 232-235. 
Azote, 117. 
Azurite, 442, 449. 

Babbitt's metal, 484. 

Baking powder, 320, 423, 438. 

Soda, 423. 
Balard, 368. 
Barite, 474. 
Barium, 473, 474, 475, 476. 

Ions, 475. 

Oxides, 17, 19, 101, 187, 474. 

Oxides, equilibrium, 187. 

Sulphate, equilibrium, 194. 
Barometer, 116. 
Baryta water, 474. 
Barytes, 474. 
Base, 152. 

And hydroxyl ions, 152. 

And metals, 151. 

And oxides, 162. 

Composition, 151. 

Diacid, 158. 

Dissociation, 164. 

Ionization, 164. 

Monacid, 158. 

Nomenclature, 161. 

See Alkali. 

Triacid, 158 
Basic, 151. 

Anhydrides, 163. 

Lining, 540. 

Oxides, 163. 

Reaction, 158. 

Salt, 158, 165. 

See Alkaline. 
Basicity, 157. 
Basil Valentine, 345. 
Bath metal, 446. 
Battery, electric, 178. 
Bauxite, 498. 



INDEX 



569 



Becher, 25. 

Beckmann apparatus, 240. 

Becquerel, 528. 

Beer, 315. 

Beet sugar, 325. 

Residues, 428, 433. 
Bell metal, 447. 
Benz aldehyde, 316. 
Benzene, 299, 329. 
Benzine, 294. 
Benzol, 329. 
Bergman, 282. 
Berlin blue, 549. 
Berthollet, 90. 
Beryl, 383, 497. 
Berzelius, 85. 
Bessemer steel, 539. 
Beta particles, 529. 
Bichromates, 523. 
Binary compounds, 162. 
Biscuit ware, 503. 
Bismite, 406. 
Bismuth, 406. 

Oxides, 407. 

Oxy chloride, 407. 

Subnitrate, 407. 

Test, 407. 
Bismuthinite, 406. 
Bisulphite, calcium, 328. 

Soda, 344. 
Bitter almonds, 316. 
Bituminous coal, 268. 
Bivalent elements, 254, 256, 257. 
Black, 282, 481. 
Black ash, 421. 

Damp, 289. 

Lead, 265, 510. 
Blast furnace, 534. 

Lamp, 38. 
Bleaching, chlorine, 200, 202. 

Hydrogen dioxide, 87. 

Sodium dioxide, 427. 

Sulphur dioxide, 342. 
Bleaching powder, 201. 
Bleach liquors, 209. 
Blooms, 538. 
Blowpipe, flame, 309. 

Mouth, 308. 

Oxyhydrogen, 37. 
Bluestone, 448. 
Blue vitriol, 448. 



Boiler scale, 473. 

Boiling point, elevation, 74, 140. 

Water, 60, 560. 
Bone ash, 391, 453. 

Cupel, 454. 
Bone black, 272. 
Bones, 391, 396. 
Boracic acid, 374. 
Boracite, 374. 
Borax, 375. 

Bead, 375, 376. 

Hydrolysis, 376. 

Reaction, 376. 

Use, 376, 473. 
Bordeaux mixture, 449. 
Boric acid, 374. 
Boride, carbon, 374 
Bornite, 442. 
Boron, 374. 

Nitride, 374. 

Oxide, 374, 376. 

Test, 375. 
Bort, 264. 
Boussingault, 119. 
Boyle's law, 46, 233. 
Brand, 391. 
Brandy, 315. 
Brass, 446. 
Braunite, 517. 
Bread, 328, 423. 
Breathing, 26. 
Brimstone, 334. 
Brin's process, 17. 
Britannia metal, 508. 
Bromides, 365, 366, 367, 368. 
Bromine, 365. 

Discovery, 368. 

Formula, 367. 

Preparation, 366. 

Properties, 367. 

Water, 367. 
Bronze, 446. 

Aluminium, 447. 

Phosphor, 446. 

Silicon, 447. 
Brucite, 480. 
Bunsen, air, 118. 

Burner, 305. 

Flame, 305. 

Spectroscope, 417, 440. 
Burette, 156. 



570 



INDEX 



Burner, acetylene, 292. 

Bunsen, 305. 

Self-lighting, 34. 
Burning, 23. 

See Combustion. 
Butter, 322. 
Butyric acid, 322. 

Cadmium, 486. 

Hydroxide, 486. 

Ions, 486. 

Sulphide, 486. 

Test, 486. 
Cesium, 417, 435, 440. 
Calamine, 482. 
Calcite, 464. 
Calcium, 463. 

Acid carbonate, 467. 

Acid phosphate, 396. 

Bicarbonate, 467. 

Bisulphite, 328. 

Borate, 375. 

Carbide, 284, 285, 290, 464-467. 

Carbonate, 464, 488. 

Carbonate, acid and normal, 281. 

Chlorate, 432. 

Chloride, 463, 471. 

Compounds, and hard water, 
473. 

Cyanamide, 472. 

Fluoride, 362, 363, 364, 365. 

Hydride, 464. 

Hydroxide, 66, 469. 

Hypochlorite, 201. 

In crust, 11, 463. 

Ions, 472. 

Light, 38. 

Magnesium carbonate, 478, 481. 

Nitrate, 217. 

Nitride, 464. 

Occurrence, 11, 463. 

Oxalate, 472. 

Oxide, 467. 

Phosphate, 391-394, 396, 400. 

Preparation, 463. 

Properties, 463, 464. 

Sulphate, 470. 

Sulphide, 421, 472. 

Sulphite, acid, 328. 

Test, 472. 
Caliche, 427. 



Calomel, 489. 
Calorie, 155, 175, 530. 
Calorific value, coal, 269. 
Calorimeter, 175. 
Candle power, 300, 309. 
Cane sugar, 324. 

See Sugar. 
Cannizzaro, 235. 
Caramel, 324. 
Carat, diamond, 265. 

Gold, 460. 
Carbide, aluminium, 289. 

Calcium, 284. 
Carbides, 275, 284. 

And petroleum, 294. 
Carbohydrates, 324. 
Carbolic acid, 329. 
Carbon, 263-275. 

Amorphous, 263, 266, 274. 

As fuel, 275. 

Atomic weight, 243. 

Bisulphide, 352. 

Boride, 374. 

Chemical properties, 274. 

Disulphide, 65, 352, 353. 

Gas, 273, 297. 

Monoxide, 282, 283, 300, 536. 

Oxides, 275. 

See Carbon dioxide. 

Silicide, 385. 

Test, 272. 

Tetrachloride, 313. 
Carbona, 313. 
Carbonado, 264. 
Carbonate, 280. 

Acid, 281. 

Normal, 281. 
Carbon dioxide, 276. 

And combustion, 276. 

Baking powder, 423. 

Composition, 281. 

Critical temperature, 278. 

Formation, 276. 

History, 282. 

In atmosphere, 121. 

Liquid, 278. 

Mines, 277. 

Occurrence, 276. 

Preparation, 277. 

Relation to life, 279. 

Solid, 278. 



INDEX 



571 



Carbon dioxide, solubility, 277, 278. 

Test, 276, 470. 
Carbonic acid, 280. 

Anhydride, 280. 

Ions, 280. 
Carborundum, 385, 496. 

Furnace, 386. 
Carboxyl, 317. 

Carnallite, 366, 428, 429, 478. 
Carnotite, 476. 
Casein, 326. 
Cassiterite, 505. 
Cast iron, 536, 537. 
Castner method, sodium, 418. 
Catalysis, 185, 288, 313. 

Chlorine, 197. 

Sulphuric acid, 349. 

Water, 200. 
Cathode, 141, 180. 
Cations, 134. 
Caustic alkali, 151. 

Lime, 468. 

Lunar, 456. 

Potash, 433. 

Soda, 424. 
Cavendish, 25, 36, 39, 82, 118, 
Caves, 465. 
Celestite, 474. 
Cell, electrolytic, 140, 179. 

Semi-permeable, 137. 

Storage, 180. 

Voltaic, 178. 
Celluloid, 329. 
Cellulose, 328. 

Nitrates, 328. 
Cement, 469, 480. 

Plaster, 471. 
Cementation process, 542. 
Cementite, 537. 
Centigrade and absolute temperature, 

43, 44. 
Cerium, 515, 516. 
Cerussite, 510. 
Chalcedony, 379. 
Chalcocite, 442. 
Chalcopyrite, 442, 532. 
Chalk, 466. 
Chalybeate water, 53. 
Chamber acid, 348. 
Change, kinds, 102. 

Physico-chemical, 5. 



Change. See Chemical and Physical 

Action. 
Chaptal, 129. 
Charcoal, 271. 

Animal, 272, 325. 

Distillation, 272. 

Filter, 271. 

Kiln, 271. 

Pit, 271. 

Wood, 271, 272, 273. 
Charles' law, 43, 233. 
Chemical action, 4, 17, 102. 

And electricity, 177. 

And heat, 171. 

And hydrogen, 32. 

And light, 169. 

And oxygen, 17, 22. 

And solution, 76, 184. 

Kinds, 102. 
Chemical compounds, 13, 78. 
Chemical energy, 7, 169, 177. 

See Energy. 
Chemical equation, 101, 102, 107. 

See Equation. 
Chemical equilibrium, 184, 186-195. 

See Equilibrium. 
Chemical equivalents, 229. 

See Equivalents. 
Chemical properties, 5. 
Chemical reaction, 6. 
Chemical symbols, 10, 11, 12, 13, 97, 
134, 564. 

And atoms, 97. 

And ions, 134. 
Chemistry, defined, 1. 

Organic, 310. 
Chile saltpeter, 426. 
China, 503. 
Chinese white, 484. 
Chlorate of potash, 432. 
Chloride, 199, 203, 206, 207, 224, 
335. 

Of lime, 201. 
Chlorination process, 459. 
Chlorine, 196-201. 

Acids, 160, 208. 

And bromides, 366. 

And iodides, 369. 

And water, 79. 

Available, 201. 

Hydrate, 199. 



572 



INDEX 



Chlorine, Ionic, 145, 208. 

Liquid, 199. 
. Oxides, 209. 

Preparation, 196, 197, 424, 425. 

Properties, 198. 

Replacing power, 367, 369. 

Water, 79, 198. 
Chloroform, 313. 
Chlorophyll, 279, 532. 
Chloroplatinic acid, 556. 
Choke damp, 289. 
Chromates, 523, 527. 
Chrome, alum, 525. 

Iron ore, 522. 

Yellow, 515, 525. 
Chromic acid, 527. 

Compounds, 526. 

Hydroxides, 526. 

Oxide, 526. 

Sulphate, 526. 
Chromite, 522, 527. 
Chromites, 526, 527. 
Chromium, 522. 

Alum, 525. 

As metal and non-metal, 525. 

Compounds, 523, 524, 525, 526. 

Ions, 527. 

Mordant, 525. 

Oxides, 526, 527. 

Preparation, 495, 522. 

Test, 524, 525. 

Valence, 526, 527. 
Chromous compounds, 526. 
Cinnabar, 486, 487, 489. 
Citric acid, 320. 
Classification of Elements, 355. ■ 

Periodic, 357, 361. 

Table, 358. 
Clay, 503. 

Ironstone, 548. 
Coal, 266. 

Anthracite, 268, 269. 

Beds, 267. 

Bituminous, 268, 269, 295. 

Brown, 269. 

Calorific value, 269. 

Classes, 268. 

Composition, 268. 

Distillation, 295. 

Distribution, 270. 

Fossil, 267. 



Coal, Gas, 295. 

Gas, composition, 300. 

Gas, plant, 295, 296. 

Hard, 269. 

Map, 270. 

Products from, 297. 

Section, 267. 

Soft, 269. 

Tar, 297. 
Cobalt, 552. 

Ions, 146. 

Test, 553. 
Cobaltite, 552. 
Coins, bronze, 446. 

Gold, 461. 

Nickel, 447, 550. 

Silver, 455. 
Coke, 273, 385. 
Colemanite, 374, 375. 
Collodion, 328. 
Colloid, 381. 
Colloidal solution, 382. 

Arsenic trisulphide, 404. 

Ferric hydroxide, 546. 

Gold, 461. 

Suspension, 382. 
Combination, 19, 22. 

And equation, 102. 
Combining capacity, 253, 257. 

Of atomic groups, 257. 

Of atoms, 257. 
Combustion, 24, 25. 

And heat, 171. 

And hydrogen, 37. 

And light, 170. 

Broad use, 35. 

Old theory, 25. 

Ordinary, 276. 

Present explanation, 25. 

Products, 304. 

Spontaneous, 24. 
Common salt. See Sodium chloride. 
Composition, air, 118. 

Ammonia gas, 216. 

Carbon dioxide, 281. 

Coal, 268. 

Earth's crust, 10, 11. 

Illuminating gas, 300. 

Of a compound, 78. 

Organic compounds, 310. 

Percentage, 1QQ, 



INDEX 



573 



Composition, Sulphur dioxide, 342. 

Water, 78-85. 
Compounds, 13, 78. 

Formula, 98, 248. 
Concentration and equilibrium, 189- 
195. 

Fraction, 194. 

Ionic, 193, 195. 

Maximum, 67. 

Of ore, 414. 

Of solutions, 63. 
Concrete, 469. 
Condenser, water, 55. 

Illuminating gas, 295. 
Conductivity, electrical, 132, 133. 

See Electrolysis. 
Condy's liquid, 520. 
Conservation, energy, 8, 169. 

Matter, 7, 169. 
Constitution, organic compounds, 
311. 

Formula, 312. 
Contact method, sulphuric acid, 348. 
Converter, 539. 
Cooking soda, 423. 
Copper, 442. 

Acetate, 450. 

Alloys, 446. 

Ammonia compounds, 450. 

And nitric acid, 219. 

And sulphuric acid, 341. 

Blister, 444. 

Carbonates, 442, 449. 

Chlorides, 447. 

Coins, 446. 

Compounds, 447. 

Displacement, 445, 450, 451. 

Electrolytic, 444. 

Family, 461. 

Fluoride, 364. 

Glance, 442. 

Ions, 146, 147, 450. 

Matte, 443. 

Metallurgy, 443. 

Nitrate, 449. 

Ore map, 533. 

Oxide, 442, 447. 

Poisonous, 447, 448. 

Properties, 445. 

Purification, 445. 

Pyrites, 442, 



Copper, Replacement, 445, 450, 451. 

See Cupric and Cuprous. 

Sulphate, 143, 444, 447, 448, 449. 

Sulphate, hydrolysis, 166, 167, 
448. 

Sulphide, 442, 449. 

Test, 450. 

Uses, 445. 
Copperas, 547. 
Coquina, 466. 
Coral, 467. 
Cordite, 323. 
Corpuscles, 96. 
Corrosive sublimate, 489. 
Corundum, 496, 497. 
Courtois, 370. 
Cracking petroleum, 300. 
Cream of tartar, 320, 423, 433. 
Critical pressure, 128. 

Acetylene, 290. 

Ammonia, 128. 

Carbon dioxide, 278. 

Ethylene, 289. 

Hydrogen, 128. 

Oxygen, 128. 
Critical temperature, 128, 290. 

Air, 128. 

Ammonia, 212. 

Carbon dioxide, 128, 278. 

Ethylene, 289. 

Nitrous oxide, 222. 

Oxygen, 128. 

Sulphur dioxide, 128, 342. 
Crockery, 504. 
Crocoisite, 522. 
Crocoite, 522. 
Crocus, 546. 

Crucible process, steel, 542. 
Cryolite, 362, 389, 493, 501. 
Crystallization, 69, 561. 

And solution, 70. 

Water of, 71, 72. 
Crystalloids, 382. 
Crystals, 69, 73. 

Systems, 561. 
Cullet, 387. 
Cupellation, 453. 
Cupric compounds, 447. 

Ferrocyanide, 450* 

Ions, 447, 448. 

Oxide, 44^ 



574 



INDEX 



Cupric compounds. See Copper. 

Sulphate, 448. 
Cuprite, 448. 
Cuprous compounds, 447. 

Ions, 447. 

Oxide, 448. 

See Copper. 
Curie, 528. 
Curve, solubility, 68. 
Cyanamide, 472. 
Cyanide, mercury, 330. 

Potassium, 434. 

Potassium auri-, 461. 

Potassium silver, 455. 

Process, 415, 459. 

Sodium, 427. 
Cyanides, 330. 

Iron, 548. 
Cyanogen, 330. 

Ions, 146. 
Cymogene, 294. 

Dalton, atomic theory, 92. 

Multiple proportions, 90. 
Davy, alkali metals, 81, 177, 198, 
222, 264, 308, 370, 417, 418, 
419, 463. 
Deacon process, 197. 
Decay, 27, 276, 288. 

Animal matter, 210, 216. 
Decomposition, 18. 

And equations, 102. 

Double, 103, 154, 207. 

Spontaneous, 530. 
Decrepitation, 430. 
Definite proportions, 89. 
Deflagration, 220. 
Degree of dissociation, 135, 163. 
Dehydrated compounds, 72. 
Deliquescence, 74, 424, 426, 432, 433, 

480. 
Denatured alcohol, 314. 
Density, 19. 

Gases, 48. 

Oxygen, 19. 

Water, 56. 
Depression of freezing point, 74, 

139. 
Destructive distillation, 289. 
Determination, atomic weights, 228, 
242-248. 



Determination, Equivalent weights, 
230. 

Molecular weights, 238, 240, 
241, 242. 
Dewar, 34, 125, 128. 
Dew-point, 120. 
Dextrose, 324, 325, 326. 
Diamond, 263. 

Allotrope, 274. 

And radium, 476. 

Artificial, 264. 

Cheap, 380. 

Crystals, 264. 

Cullinan, 265. 

Gems, 264, 265. 

Royal, 265. 
Diaspore, 498. 
Diastase, 315, 327. 
Diatomaceous earth, 379. 
Dibasic acid, 157. 
Dicalcium phosphate, 401. 
Dichromates, 523, 527. 
Diffusion, 34, 122. 
Disilicate, 383. 
Disinfectant, 316, 330, 520. 
Disodium phosphate, 396. 
Displacement, downward, 198. 

Of equilibrium, 191-195. 

Of metals, 450. 

Upward, 211. 
Dissociation, acids, 163. 

Ammonia, 212. 

Bases, 163. 

Degree, 135. 

Electrolytic, 133. 

In solution, 134, 163. 

Salts, 163. 

Table, 164. 

Water, 61, 62. 
Distillation, 54. 

Coal, 295. 

Destructive, 289. 

Dry, 210, 317. 

Petroleum, 294. 

Water, 54. 

Wood, 272, 317. 
Disulphuric acid, 352. 
Dolomite, 478, 481. 
Double decomposition, 103, 207. 

Neutralization, 154. 
Double salt, 428, 501. 



INDEX 



575 



Downward displacement, 198. 
Drinking water, 53. 

And ozone, 29. 
Dulong and Petit, 245. 
Dumas, 84, 119, 264, 356. 
Dutch metal, 446. 

Process, white lead, 514. 
Dyads, 257. 
Dyeing, 501. 
Dynamite, 323. 

Earth's crust, 11. 
Effervescence, 278. 
Efflorescence, 72, 422, 426. 
Electric charges, 134, 142. 

Current, 179. 

Furnace, 173, 285, 385. 
Electrical energy, measurement, 183. 
Electricity and chemical change, 177. 

And solution, 132. 
Electrochemical equivalent, 183. 

Series, metals, 451. 
Electrochemistry, 182. 
Electrodes, 141, 179, 266. 
Electrolysis, 179, 183. 

Aluminium, 493. 

Applications, 181. 

Calcium chloride, 463. 

Carnallite, 478. 

Copper sulphate, 143, 444. 

Extraction by, 415. 

Fused salts, 144, 417, 478, 493. 

Gold, 460. 

Hydrochloric acid, 141. 

Illustrations, 143. 

Interpretation, 180. 

Rusting, 544. 

Silver, 455. 

Sodium chloride, 424. 

Sodium hydroxide, 417. 

Sodium sulphate, 143. 

Solution, 140. 

Theory, 140-144, 180. 

Water, 80, 143. 
Electrolytes, 133, 136. 
Electrolytic cell, 140, 179, 455. 
Electrolytic dissociation, 133, 544. 
Electrolytic process, chlorine, 198. 

Sodium hydroxide, 417. 
Electrolytic solution, 133. 

Boiling point, 140. 



Chemical behavior, 144. 

Freezing point, 139. 
Electron, 96, 529, 531. 
Electro-negative ions, 134, 146, 147. 
Electroplating, 182. 
Electro-positive ions, 134, 146, 147. 
Electrosilicon, 380. 
Electrotyping, 181, 182. 
Elements, 8-13, 358, 564. 

Acid-forming, 150, 413. 

Base-forming, 151, 413. 

Bivalent, 254, 256, 257. 

Classification, 355. 

Decay, 530, 531. 

Families, 356, 359. 

Groups, 356, 359. 

In crust, 11. 

Molecule of, 95, 250. 

Numerical relations, 356. 

Periodic classification, 357. 

Periodic table, 358. 

Properties, acid, 150, 355, 413. 

Properties, basic, 151, 355, 413. 

Quadrivalent, 254, 255, 257. 

Quinquivalent, 257. 

Radioactive, 476, 528-531. 

Tables, 10, 11, 12, 564. 

Trivalent, 254, 255, 257. 

Univalent, 254-257. 

Valence, 254-257. 
Elevation of boiling point, 74, 140. 
Emerald, 497. 
Emery, 496. 
Empirical formula, 312. 
Endothermic compounds, 176, 292, 

352. 
Energy, 7, 8, 274. 

And light, 170. 

Chemical, 7, 169, 177. 

Conservation, 8. 

Electrical, 177, 183. 

Heat, measurement, 175. 

See Electrolysis. 

Storage cell, 180. 
Enzymes, 315, 325, 327, 328. 
Epsom salts, 480. 
Equation, 101. 

And calculations, 111. 

Gas, 252. 

Ionic, 154. 

Making, 107. 



576 



INDEX 



Equation, Molecular, 251. 

Neutralization, 153, 156. 
. Ordinary, 102. 

Preliminary, 18. 

Reversible, 189. 

Thermal. 176. 

Volume, 252. 
Equilibrium, 61, 70, 184, 186-195. 

Ammonia, 212. 

And solution, 189, 193-195. 

Barium oxides, 187. 

Barium sulphate, 194. 

Concentration, 191-195. 

Constant, 191. 

Displacement, 191-195. 

Heat, 191, 192, 224. 

Hydriodic acid, 189, 190, 371. 

Hydrochloric acid, 204. 

Hydrogen, 187-191. 

Le Chatelier's law, 191, 224. 

Nitric acid, 217. 

Nitrogen oxides, 224. 

Phosphorus pentachloride, 398. 

Saturated solution, 70. 

Steam and iron, 187, 188, 191. 

Water vapor, 61. 
Equivalent weights, 229. 

And atomic weights, 231, 244. 

And valence, 254. 

Determination, 230. 

Table, 230. 
Equivalents, chemical, 183. 

Electrochemical, 183. 

See Equivalent weights. 
Erosion, 51. 
Esters, 320. 
Etching, 365. 
Ether, ethyl, 316. 

Solubility in water, 66. 

Sulphuric, 316. 
Ethyl, 311, 312. 

Acetate, 320. 

Alcohol, 312, 314. 

Butyrate, 321. 

Ether, 316. 
Ethylene, 289. 

Evaporation of water, 57-61. 
Exothermic compounds, 176. 
Explosions, acetylene, 290. 

Coal mine, 289. 

Hydrogen, 35. 



Factors, 106. 

Fahrenheit thermometer, 560. 

Faraday, 127, 179, 183, 199. 

Law, 183. 
Fats, 321, 322. 
Fehling's solution, 326. 
Feldspar, 503. 
Ferment, 314, 318. 
Fermentation, 315. 

Acetic, 318. 

Alcoholic, 277, 315. 

Bread, 327. 

Sugar, 325, 326. 
Ferric compounds, 544. 

Chloride, 548. 

Ferrocyanide, 549. 

Hydroxide, 546. 

Oxide, 546. 

Sulphate, 547. 

Sulphide, 547. 

Sulphocyanate, 549. 
Ferricyanides, 548, 549. 
Ferrite, 546. 
Ferro chrome, 543. 
Ferrocyanides, 450, 548, 549. 
Ferromanganese, 517, 540. 
Ferrosilicon, 543. 
Ferroso-ferric oxide, 546. 
Ferro tungsten, 543. 
Ferrous compounds, 544. 

Carbonate, 548. 

Chloride, 548. 

Chromite, 522. 

Ferric oxide, 546. 

Ferri cyanide, 549. 

Ferrocyanide, 549. 

Hydroxide, 546. 

Oxide, 546. 

Sulphate, 222, 546. 

Sulphide, 547. 
Fertilizer, 131, 400, 401, 426, 434, 

437, 541. 
Filter, charcoal, 271. 
Fire, 24. 

And carbon oxides, 282. 

Damp, 288. 

Extinguisher, 278. 

See Combustion. 
Fireworks, 24, 474. 
Fixation, nitrogen, 131. 
Fixed air, 282. 



INDEX 



577 



Fixed alkali, 161. 
Flame, 301. 

Acetylene, 290, 291. 

And gauze, 307. 

Bunsen, 305, 306, 307. 

Candle, 302, 303. 

Hydrogen, 36. 

Lamp, 302. 

Luminosity, 304. 

Luminous, 302, 305. 

Non-luminous, 305, 306. 

Ordinary gas, 304. 

Oxidizing, 308. 

Oxy acetylene, 27. 

Oxyhydrogen, 37. 

Parts, 302, 303. 

Structure, 302, 303. 

Reducing, 308. 
Flashing point, 294. 
Flint, 379. 

Flowers of sulphur, 334. 
Fluid magnesia, 481. 
Fluorides, 365. 
Fluorine, 362. 

Compounds, 362. 

Formula, 363. 

Isolation, 362. 
Fluorite, 362. 
Fluor spar, 362. 
Fluosilicic acid, 384. 
Flux, 415, 535. 
Fool's gold, 547. 
Formaldehyde, 316. 
Formalin, 316. 
Formula, 97. 

And valence, 258. 

Calculation, 101. 

Constitutional, 260. 

Determination, 248. 

Empirical, 312. 

Graphic, 260, 312, 317. 

Molecular, 248. 

Rational, 312. 

Simplest, 249, 250. 

Structural, 260, 312, 317. 

Weight, 238. 

Writing, 258. 
Forward reaction, 188, 190, 192. 
Franklinite, 482. 
Freezing point, 560. 

And molecular weight, 239. 



Freezing point, And solution, 139. 

Depression, 139. 

Determination, 240, 241, 248. 

Solution. 74, 
Fructose, 326. 
Fruit sugar, 326. 
Fuming acid, nitric, 223. 

Sulphuric, 351. 
Furnace, blast, 534. 

Electric, 173, 174. 

Open hearth, 541. 

Reverberatory, 415. 
Fusible alloys, 407. 

Metals, 407, 486. 

Gahnite, 482. 

Galena, 510, 514. 

Galvanized iron, 483. 

Gangue, 414. 

Garnet, 497. 

Garnierite, 550. 

Gas, burner, self-lighting, 34. 

Carbon, 273, 297. 

Coal, 295. 

Equation, 252. 

Flame, 302, 303. 

Illuminating, 295. 

Liquor, 211. 

Marsh, 288. 

Natural, 288, 295. 

Olefiant, 289. 

Pintsch, 299. 

Producer, 32, 283. 

Reaction, 189. 

Water, 299. 
Gases, adsorption by charcoal, 271. 

And pressure, 41, 42, 46. 

And temperature, 41, 42-46. 

Density, 48. 

Gay-Lussac's law, 224, 233. 

Inert, 123. 

Kinetic theory, 232. 

Liquefaction, 127. 

Molecular formula, 250. 

Solution, 64. 
Gasolene, 294. 
Gastric juice, 204. 
Gay-Lussac, 83, 346, 370. 

Law, 224, 233. 

Tower, 346. 
Gems, aluminium, 497. 



578 



INDEX 



Gems, Artificial, 497. 
German silver, 447. 
Geyserite, 384. 
Glacial acetic acid, 318. 

Phosphoric acid, 395. 
Glass, 387. 

Annealing, 388. 

Blowing, 389. 

Colored, 389. 

Crown, 389. 

Cut, 388. 

Etching, 365. 

Flint, 388. 

Hard, 389. 

Kinds, 388. 

Lead, 388. 

Manufacture, 388. 

Quartz, 380. 

Water, 383. 

Window, 389. 
Glauber salt, 426. 
Glazing, 503, 504. 
Globigerina ooze, 466. 
Glover tower, 346. 
Glucose, 326. 
Glycerin, 321, 322. 
Glycerol, 323. 
Glyceryl, 321. 

Oleate, 321. 

Palmitate, 321. 

Stearate, 321. 
Gold, 457. 

Alloys, 458, 461. 

Amalgam, 488. 

Carat, 460. 

Coin, 461. 

Colloidal, 461. 

Compounds, 461. 

Cyanide process, 459. 

Electrolytic purification, 460. 

Fool's, 547. 

History, 458. 

Leaf, 460. 

Metallurgy, 458. 

Mining, 458. 

Occurrence, 458. 

Ores, 458. 

Plating, 461. 

Purification, 459. 

Purity, 458, 460. 

Red, 461. 



Gold, Separation from silver, 459. 

Test, 461. 

Uses, 461. 

White, 461. 
Goldschmidt method, 496, 497. 
Graham, 34, 382. 
Gram, 557. 

Gram-molecular volume, 238. 
Gram-molecular weight, 138, 189, 

240. 
Grape sugar, 326. 
Graphic formula, 260, 291, 312. 
Graphite, 265, 266. 

Allotrope, 274. 

Artificial, 266. 
Gravimetric composition, air, 118. 

Water, 83. 
Green fire, 474. 

Vitriol, 546. 
Greenockite, 486. 
Guano, 400. 
Guignet's green, 526. 
Gun cotton, 328. 

Metal, 447. 
Gunpowder, 75, 271, 431. 
Gypsum, 470. 

Hsemoglobin, 283, 532. 
Halides, 362. 
Halogens, 362. 

Acids, ions, 372. 

Family, 372. 
Hardness, water, 473. 
Hausmannite, 517. 
Heat and chemical change, 171. 

Combustion, 176. 

Decomposition, 176. 

Electric furnace, 173. 

Equilibrium, 191, 192, 224. 

Formation, 176. 

Le Chatelier's law, 191, 224. 

Measurement, 175. 

Of neutralization, 155. 

Of solution, 75. 

Sources, 173. 

Velocity reaction, 191, 224. 
Helium, 123. 

And radium, 123, 531. 

Detection, 441. 

Spectrum, 124. 
Hematite, 532, 533, 534. 



INDEX 



579 



Henry's law, 65. 
Heptads, 257. 
Heptavalent elements, 257. 
Heterogeneous mixture, 189. 
Hexads, 257. 

Hexavalent elements, 257. 
Homogeneous mixture, 189, 387. 
Humboldt, 83. 
Humidity, 120. 
Hydrargyllite, 498. 
Hydrate, 71. 

Chlorine, 199. 
Hydraulic main, 296. 
Hydriodic acid, equilibrium, 189, 190, 

371. 
Hydrobromic acid, 368. 
Hydrocarbons, 288, 299, 313. 
Hydrochloric acid, 203, 205, 421. 

Commercial, 204. 

Electrolysis, 141. 

Equilibrium, 204. 

Ionic test, 145. 

Ions, 206. 
Hydrocyanic acid, 330. 
Hydrofluoric acid, 364, 379. 

Electrolysis, 363. 
Hydrofluo silicic acid, 384. 
Hydrogen, 30-39. 

And acids, 31, 150, 452. 

And alpha particles, 529. 

And chlorine, 35, 199, 205. 

And iodine, 189, 190. 

And metals, 452. 

And oxygen explosion, 35. 

And water, 31, 33, 78. 

Arsenide, 404. 

As metal, 452. 

Atom, 236, 529. 

Bromide, 368. 

Chloride, 203, 236. 

Diffusion, 34. 

Dioxide, 86, 186, 248. 

Displacement, 452. 

Equilibrium, 187-191. 

Explosion, 33, 35. 

Flame, 36, 175. 

Fluoride, 364. 

From nitric acid, 219. 

Ions, 146, 150, 452. 

Lavoisier's experiment, 187. 

Liquid, 128. 



Hydrogen, Manufacture, 425. 

Molecule, 236. 

Molecular equation, 252. 

Peroxide. See Dioxide. 

Solid, 128. 

Standard, 228, 236. 

Weight of liter, 38. 
Hydrogen sulphide, 337, 339. 

Composition, 339. 

Ions, 339. 

Test, 340. 
Hydrolysis, 166. 

Acid sodium sulphate, 166. 

Aluminium compounds, 499, 501, 
502. 

Antimony trichloride, 406. 

Bismuth trichloride, 407. 

Copper sulphate, 165, 448. • 

Ferric chloride, 166, 548. 

Magnesium chloride, 480. 

Phosphates, 397. 

Potassium carbonate, 166. 

Potassium cyanide, 166. 

Soap, 324. 

Sodium carbonate, 165, 422. 

Sugar, 325, 326, 327. 

Sulphites, 345. 

Zinc chloride, 485. 
Hydroquinone, 457. 
Hydroxyl, 147, 152, 154. 
Hydroxyl ions, bases, 152. 
Hypo, 352, 457. 
Hypochlorous acid, 201, 208. 
Hypophosphites, 397. 
Hyposulphite, sodium, 352. 
Hypothesis, 88. 

Avogadro's, 232-235. 

Ice, 57. 

Plant, 215. 
Iceland spar, 464. 
Iiluminants, 299, 300. 
Illuminating gas, 295, 299. 

Candle power, 300, 301. 

Characteristics, 300. 

Composition, 300. 

Impurities, 300. 

Luminosity, 300. 

Tarnishing by, 454. 
Inert gases, 123, 124. 
Infusorial earth, 379. 



580 



INDEX 



Ingots, 542. 

Ink, 456, 547. 

Inorganic compounds, 310. 

Insecticides, 337, 403. 

Insoluble substances, 63. 

Invertase, 325. 

Iodides, 368, 369, 371, 373. . 

Iodine, 189, 190, 368. 

And chlorine, 371. 

And hydrogen, 189, 190. 

And starch, 185, 327. 

Preparation, 369. 

Properties, 370. 

Purification, 369, 370. 

Saltpeter, 368, 369, 427. 

Seaweed, 368, 369. 

Solution, 372. 

Test, 370. 

Vapor density, 370. 
Iodoform, 312, 372. 
Ionic equation, 154. 

Concentration, 193, 195. 

Hydrolysis, 167. 

Neutralization, 195. 

Test, sulphate, 351. 
Ionization, 134, 193-195. 

And salts, 165, 193-195. 

See Dissociation, Electrolytic. 
Ions, defined, 134. 

And atoms, 134. 

And electrolysis, 180. 

And equilibrium, 187, 188, 191, 
193-195. 

And neutralization, 154. 

And water, 165, 194, 422, 448, 
449. 

Color, 146. 

Common, 146. 

Complex, 146. 

Effect of removal, 191. 

Interaction, 144. 

Migration, 141, 146, 180, 424, 
425. 

Symbols, 134. 

Table, 147. 

Water, 165, 194, 422, 448, 449. 
Iridium, 554, 555, 556. 
Iridosmine, 556. 
Iron, 532. 

Acetate, 319. 

Alum, 547. 



Iron, And steam, 187, 188, 190. 
Cast, 536, 537. 
Chemistry of smelting, 536. 
Chlorides, 548. 
Compounds, 545. 
Cyanides, 548. 
Disulphide, 547. 
Electrolytic preparation, 544. 
Equilibrium, 187, 188, 190. ' 
Ferrite, 546. 
Flux, 535. 
Galvanized, 483. 
Hydroxides, 545. 
Impurities, 536, 537. 
In crust, 11, 532. 
Ions, 550. 
Liquor, 319. 
Magnetic oxide, 546. 
Metallurgy, 534. 
Ore, 532, 533. 
Ore, chrome, 522. 
Ore, map, 533. 
Oxides, 297, 545. 
Passive, 544 . 
Pig, 536. 
Pure, 544. 
Pyrites, 532, 547. 
Russia, 546. 
Smelting, 534. 
Spiegel, 517. 
Steel, 538. 
Sulphate, 546. 
Sulphides, 547. 
Terms, 532. 
Test, 549. 
Valence, 545. 
Varieties, 536. 
Wrought, 537. 

Jasper, 379. 
Javelle's water, 209. 

Kainite, 428, 434, 478. 

Kaolin, 383, 503. 

Kerosene, 294. 

Kieserite, 478, 480. 

Kindling temperature, 25, 172, 304, 

307. 
Kinetic theory, 232. 
Kirchhoff, 440. 
Krypton, 123. 



INDEX 



581 



Labarraque's solution, 209. 

Lactic acid, 423. 

Lactose, 326. 

Lake, 501. 

Lampblack, 274. 

Laughing gas, 222. 

Lavoisier, 25, 27, 31, 37, 39, 83, 117, 

163, 264, 282. 
Law, 88. 

And theory, 88. 

Boyle, 46, 233. 

Charles, 43, 44, 45, 233. 

Conservation of energy, 8. 

Conservation of matter, 7, 93. 

Definite proportions, 89, 94. 

Dulong and Petit, 245. 

Faraday, 183. 

Gay-Lussac, 224, 233. 

Henry, 65. 

Le Chatelier, 191, 224. 

Mass action, 189-195. 

Multiple proportions, 90, 94. 

Periodic, 361. 

Specific heat, 245. 
Lead, 510. 

Acetate, 319, 515. 

Alloys, 512. 

Argentiferous, 453. 

Arsenate, 403. 

Black, 265, 510. 

Carbonate, 510, 513. 

Chloride, 515. 

Chromate, 515, 522, 525. 

Cupellation, 453. 

Dioxide, 513. 

Displacement, 511. 

Fluo silicate, 511. 

Glass, 388. 

In drinking water, 512. 

Ions, 515. 

Metallurgy, 510. 

Monoxide, 513. 

Nitrate, 515. 

Oxides, 513. 

Pencils, 266. 

Peroxide, 513. 

Phosphate, 510. 

Poisoning, 512. 

Properties, 511. 

Purification, 511. 

Red, 513. 



Lead, Storage cell, 180. 

Sugar of, 319. 

Sulphate, 510, 515. 

Sulphide, 510, 514. 

Test, 514, 515. 

Tetroxide, 513. 

White, 513, 514. 
Leblanc process, 421, 433. 
Le Chatelier's law, 191, 224. 
Leguminous plants and nitrogen, 131. 
Levulose, 326. 
Liebig, 173, 368. 
Life and nitrogen, 130. 

And oxygen, 26. 

And phosphorus, 400. 

And potassium, 434. 
Light and chemical change, 169. 

And energy, 170. 

And silver salts, 456. 
Lignite, 268, 269. 
Lime, 467. 

And bleaching powder, 201. 

Caustic, 468. 

Chloride of, 201. 

Light, 38, 467. 

Milk of, 470. 

Preparation, 468. 

Quick, 468. 

See Calcium oxide. 

Slaking, 171, 467. 

Soda, 469. 

Sulphur, 337. 

Superphosphate, 400. 

Uses, 468. 
Limekiln, 468, 469. 
Limestone, 464, 468. 
Limewater, 469. 

See Calcium hydroxide. 
Limonite, 532, 533. 
Link, valence, 312. 

Fusible, 406. 
Liquefaction of gases, 127. 
Liquefied ammonia, 212, 215. 
Liquid air, 17, 125, 127. 

Acetylene, 290. 

Ammonia, 212, 215. 

Carbon dioxide, 278. 

Chlorine, 199. 

Ethylene, 289. 

Hydrogen, 34, 128. 

Oxygen, 20, 126. 



582 



INDEX 



Liquid, Ozone, 28. 

Sulphur dioxide, 342. 
Liquids, solubility, 65. 
Liquor, alcoholic, 315. 

Distilled, 315. 

Iron, 319. 

Red, 319, 501. 
Litharge, 453, 454, 513. 
Lithium, 417, 435. 
Lithophone, 474. 
Litmus, acids, 150. 

Bases, 151. 

Neutral, 152. 

Salts, 152, 165. 
Loadstone, 546. 
Lockyer, 123. 
Lubricating oils, 294. 
Luminosity, 300, 302. 
Luminous paint, 474. 
Lunar caustic, 456. 
Luster, 412. 
Lye, 151. 

Magnalium, 496. 
Magnesia, 480, 481. 

Alba, 481, 518. 

Black, 518. 

Fluid, 481. 

Mixture, 480. 

See Magnesium. 

Stone, 518. 
Magnesite, 480, 481. 
Magnesium, 478. 

Alloy, 496. 

Aluminate, 497. 

Ammonium phosphate, 481. . 

Basic carbonate, 481. 

Bromide, 365, 366, 367. 

Calcium carbonate, 478, 481. 

Carbonate, 481. 

Chloride, 480. 

Citrate, 481. 

Compounds and water, 473. 

Electrolytic preparation, 478. 

Hydroxide, 480. 

Ions, 481. 

Nitride, 129, 479. 

Oxide, 480. 

Phosphate, 481. 

Properties, 479. 

Stassfurt salts, 478. 



Magnesium, Sulphate, 426, 480. 

Test, 481. 
Magnetic oxide of iron, 546. 
Magnetite, 532, 533. 
Malachite, 442, 449. 
Malt, 315. 
Maltose, 315, 327. 
Manganates, 520. 
Manganese, 517. 

Alloys, 517. 

Black oxide, 518. 

Compounds, 518, 519, 520. 

Dioxide, 196, 517, 518. 

Ferro-, 517. 

Ions, 520. 

Oxides, 517. 

Test, 518, 520. 

Valence, 520. 

Weldon process, 197, 518. 
Manganic acid, 520. 
Manganite, 517. 
Manganous chloride, 518, 520. 

Hydroxide, 518. 

Sulphate, 520. 

Sulphide, 520. 
Mantle, Welsbach, 309, 516. 
Marble, 464. 
Marchand tube, 83. 
Marsh gas, 288. 
Marsh's test for arsenic, 404. 
Mass action, 189-195. 
Massicot, 513. 
Matches, 399. 
Matter, conservation, 7. 

Defined, 6. 

Distribution, 11. 

Properties, 1. 
Mendelejeff, 357, 361. 
Mercuric compounds, 488. 

Chloride, 489, 509. 

Cyanide, 330. 

Ions, 147, 490. 

Nitrate, 490. 

Oxide, 15, 18, 23, 191, 488. 

Sulphide, 486, 487, 489. 
Mercurous compounds, 488. 

Chloride, 489, 509. 

Ions, 147, 490. 

Nitrate, 490. 
Mercury, 486. 

Amalgams, 488. 



INDEX 



583 



Mercury, Ammono compounds, 491. 

Compounds, 488. 

Fulminate, 352. 

Preparation, 487. 

Properties, 487. 

See Mercuric and Mercurous. 

Specific heat, 245. 

Test, 490. 
Metaboric acid, 375. 

Phosphoric acid, 395. 

Silicate, 383. 

Silicic acid, 381, 383. 

Stannic acid, 507. 
Metal and non-metal, 355, 410, 
411. 

Babbitt's, 484. 

Bath, 446. 

Bell, 447. 

Britannia, 508. 

Dutch, 446. 

Gun, 447. 

Hypothetical, 417. 

Muntz, 446. 

Newton's, 407. 

Noble, 224, 452, 458, 554. 

Rose's, 407. 

Speculum, 447. 

Type, 512. 

White, 508. 

Wood's, 407, 486. 
Metallic ions, 134, 135, 147, 413. 

See Ions and Electrolysis. 
Metalloids, 410. 
Metallurgy, 414. 

See Individual metals. 
Metals, alkali, 417, 435. 

Alkaline earth, 463, 475. 

Ancient, 413. 

And bases, 151, 413. 

And non-metals, 355, 410, 411. 

And steel, 543. 

Atomic state, 451. 

Chemical properties, 413. 

Classification, 411. 

Displacement, 450. 

Electrochemical series, 451. 

Electromotive series, 451. 

Electropositive, 451, 452. 

Fusible, 407, 486. 

Ionic state, 451. 

Ions, 134, 135, 147, 413. 



Metals, Latin names, 13. 

Native, 413. 

Occurrence, 413. 

Physical properties, 412. 

Platinum, 554, 556. 

Preliminary treatment, 414. 

Preparation, 414. 

Table, 411. 
Metathesis, 103. 
Meteorites, 532, 550. 
Meter, 557. 
Methane, 288, 308. 

In natural gas, 295. 

Series, 293. 
Methyl, 311, 312. 

Alcohol, 240, 314. 

Salicylate, 321. 
Metric system, 557. 

Equivalents, 559. 

Table, 559. 

Transformations, 559. 
Mexican onyx, 465. 
Meyer, Victor, 239. 
Microcosmic salt, 396. 
Migration of ions, 142, 180. 

See Electrolysis. 
Milk of lime, 470. 

Sour, 423. 

Sulphur, 337. 
Mineral, defined, 413. 

Orange, 513. 

Springs, 52. 

Water, 52. 
Minium, 513. 
Mispickel, 401. 
Mixture, 14, 90. 

Air, 124. 

Bordeaux, 449. 

Heterogeneous, 189. 

Homogeneous, 189, 387. 
Moissan, 264, 285, 363, 364, 463. 
Molar weight, 238. 
Molasses, 325. 
Mole, 138, 176, 189, 240. 
Molecular depression, 240. 

Equation, 251. 

Formula, 248, 250. 
Molecular weight, 99, 236, 250. 

And vapor density, 236. 

Approximate, 241. 

Depression, 240. 



584 



INDEX 



Molecular weight, Determination, 
238, 240, 241, 242. 

Exact, 241. 

Freezing point, 239. 

Gram, 138, 189, 240. 

Hydrogen dioxide, 248. 

Hydrogen standard, 236. 

Oxygen standard, 236. 

Preliminary, 99. 
Molecules, 92, 95. 

And atoms, 95. 

And dissociation, 133, 135. 

And equations, 251, 252. 

And equilibrium, 186-189, 192, 
212. 

And formulas, 97, 248-252. 

And kinetic theory, 233. 

See Atoms in molecule. 

Undissociated, 135. 
Molybdenum, 527. 
Monacid base, 158. 
Monads, 257. 
Monazite, 516. 
Monde process, 551. 
Monel metal, 551. 
Monobasic acid, 157. 
Mono clinic crystal, sulphur, 336. 
Mordant, 501, 509, 525, 547. 
Morley, 84. 
Mortar, 470. 
Moth balls, 329. 
Multiple proportions, 90, 91. 
Muntz metal, 446. 
Muriate of ammonia, 436. 
Muriatic acid, 203. 

Names, acids, 160. 

Bases, 161. 

Salts, 161. 
Naphtha, 294. 
Naphthalene, 329. 
Natural gas, 295. 
Natural groups, 356, 357, 358, 359, 

360. 
Natural waters, 52. 
Negative electrode, 141, 180. 

Photographic, 457. 
Neon, 123. 

Neutral reaction, 152. 
Neutralization, 153. 

And ionization, 154, 194. 



Neutralization, Equation r 177. 

Heat of, 155. 
Newton's metal, 407. 
Niccolite, 550. 

Nicholson and Carlisle, 81, 177. 
Nickel, 550. 

Ammonium sulphate, 550, 551. 

Carbonyl, 551. 

Chloride, 551. 

Coins, 447, 550. 

Compounds, 551. 

Hydroxide, 551. 

Ions, 146, 552. 

Plating, 550. 

Steel, 551. 

Sulphate, 551. 

Sulphide, 551. 

Test, 551. 
Nickeloid, 551. 
Niton, 358, 359, 530, 531. 
Nitrates, 219. 

Behavior with heat, 220. 

Formation, 216. 

Test, 220. 
Nitric acid, 216. 

And iron, 544. 

And metals, 219. 

Equilibrium, 217. 

Formation, 216. 

Fuming, 223. 

Ions, 218. 

Manufacture, 218. 

Preparation, 217. 

Properties, 218. 

Test, 220. 
Nitric oxide, 222. 
Nitride, magnesium, 129, 479. 
Nitrification, 216. 
Nitrites, 220, 221. 
Nitrobenzene, 329. 
Nitrogen, 129, 131. 

Additional properties, 129. 

And life, 130. 

Assimilation, 131. 

Dioxide, 223, 224. 

Family, 408. 

Fertilizer, 437. 

Fixation, 131. 

General properties, 117. 

In atmosphere, 117. 

Liquid, 129. 



INDEX 



585 



Nitrogen, Oxides, 219, 221. 

Pentoxide, 224. 

Peroxide, 223, 224. 

Tetroxide, 224. 

Trioxide, 224. , 
Nitroglycerin, 323. 
Nitrosylsulphuric acid, 346. 
Nitrous acid, 221. 
Nitrous oxide, 221, 438. 
Noble metal, 452, 458, 554. 
Nomenclature, 161. 

Acids, 160. 

Bases, 161. 

Ions, 134, 146. 

Salts, 157, 161, 162. 
Non-electrolytes, 133, 136. 
Non-electrolytic solution, 143. 

Freezing point, 139. 
Non-luminous flame, 305. 
Non-metallic ions, 134, 146, 147. 

See Ions. 
Non-metals, 150, 355, 410, 411. 

And acids, 150. 

And metals, 355, 410. 

Table, 411. 
Nordhausen sulphuric acid, 351. 
Normal salts, 157, 165. 

Solution, 163. 

Occlusion, 34. 
Ocean water, 53. 

Chlorine in, 196. 

Composition, 11. 

Solids, 53. 

Table, 11. 
Oil, lubricating, 294. 

Natural, 322. 

Of vitriol, 345. 
Olefiant gas, 289. 
Olein, 321, 322. 
Oleomargarine, 322. 
Olivine, 383. 
Onyx, 379, 465. 
Opal, 379. 
Open-hearth furnace, 541. 

Process, 541. 
Orange mineral, 513. 
Ordinary chemical equation, 102. 
Ore, defined, 413. 

Concentration, 414. 

Dressing, 414. 



Ore, Smelting, 414. 
Organic acids, 319. 

Compounds, 310, 311, 312. 
Orpiment, 401, 403. 
Orthoclase, 383. 
Orthophosphoric acid, 395. 
Orthorhombic crystals, sulphur, 336. 
Orthosilicic acid, 381, 383. 
Osmiridium, 556. 
Osmium, 556. 
Osmotic pressure, 136. 
Oxalic acid, 319. 

And carbon monoxide, 283. 
Oxidation, 22. 

And decay, 26. 

And life, 26. 

And ozone, 28. 

And reduction, 39. 

Broad meaning, 490, 509, 545. 

Ionic standpoint, 545. 

Nitric acid, 218. 

Potassium dichromate, 524. 

Potassium permanganate, 519. 
Oxide, defined, 22. 

Carbonic, 280. 
Oxides and acids, 162, 413. 

And bases, 162, 413. 

Chlorine, 209. 

Names, 24. 

Nitrogen, table, 221. 
Oxidizing agent, 24. 

Hydrogen dioxide, 87. 

Nitrate, 220. 

Nitric acid, 218. 

Potassium chlorate, 24. 

Potassium dichromate, 524. 

Potassium permanganate, 519. 
Oxidizing flame, 308. 
Oxone, 17, 427. 
Oxygen, 15-22, 26, 27. 

And acids, 27, 160. 

And blood, 26. 

And combustion, 24. 

And ozone, 28, 29. 

And water, 79. 

Breathing pure, 26. 

Brin's process, 17. 

Discovery, 27. 

Free, 208. 

In atmosphere, 117. 

Liquid, 20. See Liquid air. 



586 



INDEX 



Oxygen, Molecule, 237. 

Preparation, 16, 17, 427. 

Properties, 19, 20. 

delation to life, 26. 

Solid, 20. 

Solubility, 65. 

Standard, 228, 236. 

Weight of liter, 19. 
Oxyhydrogen blowpipe, 37. 
Oxymuriate of tin, 509. 
Ozone, 28, 29. 

Formula, 251. 

Paint, 474, 484, 514. 

Pakfong, 447. 

Palladium, 34, 556. 

Palmitic acid, 321. 

Palmitin, 322. 

Palm oil, 322. 

Paper, 328. 

Paracelsus, 39. 

Paraffin, 294. 

Paris green, 319, 403, 450. 

Paris, plaster of, 471. 

Parkes process, 453. 

Partial pressure, 60, 65. 

Passive iron, 544. 

Paste, glass, 388. 

Pearlash, 433. 

Pentad, 257. 

Percentage composition, 100. 

Periclase, 480. 

Periodic classification, 361. 

Law, 361. 
Permanent hardness, 473. 
Peroxide, hydrogen, 86, 186, 249. 

Sodium, 427. 
Petrified wood, 379. 
Petroleum, 293. 

Cracking, 300. 

Distillation, 294. 

Products, 294. 

Refining, 294. 

Well, 293. 
Pewter, 508. 
Phenol, 329. 
Phenyl, 311. 

Philosopher's stone, 458. 
Phlogiston, 25, 27. 
Phosgene, 284. 
Phosphates, 395, 396. 



Phosphates, Acid calcium, 396. 

Dicalcium, 401. 

Disodium, 396. 

Ions, 397. 

Meta-, 396. 

Primary, 396. 

Rock, 400. 

Secondary, 396. 

Slag, 400. 

Sodium ammonium, 396, 541. 

Tertiary, 396. 

Tests, 396. 

Tricalcium, 391-394, 396. 
Phosphine, 397. 
Phosphonium compounds, 398. 
Phosphor bronze, 446. 
Phosphorescence, 476. 
Phosphoric acids, 391, 392, 395. 

Glacial, 395. 

Ions, 397. 
Phosphoric oxides, 394. 
Phosphorite, 391. 
Phosphorous oxide, 394. 
Phosphorus, 391. 

Acids, 391, 392, 395. 

And nitrogen, 129. 

And ozone, 28. 

Black, 394. 

Bronze, 446. 

Burns, 393. 

Electrolytic manufacture, 392. 

Formula, 394. 

Oxides, 394. 

Pentachloride, 398, 399. 

Preparation, 391, 392. 

Properties, 393. 

Red, 394. 

Relation to life, 400. 

Sulphide, 399. 

Test, 397. 

Trichloride, 398. 

Yellow, 393. 
Photography, 170, 328, 352, 368, 

456. 
Phylloxera, 337. 
Physical change, 3. 

Properties, 4. 
Physico-chemical change, 5. 
Picromerite, 428. 
Pig iron, 536. 
Pinchbeck, 446. 



INDEX 



587 



Pintsch gas, 299. 
Pitchblende, 476, 528. 
Plaster, 470. 

Of Paris, 471. 
Platina, 554. 
Platinum, 554. 

Alloys, 555, 556. 

Arsenide, 554. 

Black, 555. 

Catalyzer, 348, 349, 555. 

Chloroplatinic acid, 556. 

Foil, 555. 

Metals, 554, 556. 

Properties, 555. 

Sponge, 554, 555. 

Sulphuric acid manufacture, 348. 

Tetrachloride, 556. 

Using," 555. 
Plumbago, 265. 
Polariscope, 326. 
Polonium, 531. 
Polyhalite, 428. 
Poly sulphides, 437. 
Porcelain, 503. 
Portland cement, 469. 
Positive electrode, 141, 180. 
Potash, 432, 433. 

Caustic, 433. 

Prussiate of, 548. 
Potassium, 428. 

And water, 429. 

Antimonyl tartrate, 406. 

Auricyanide, 461. 

Bicarbonate, 432. 

Bichromate, 523, 524. 

Bromide, 368. 

Carbonate, 432, 433. 

Chlorate, 16, 18, 429, 430, 431. 

Chloride, 432. 

Chloroplatinate, 556. 

Chromate, 523. 

Cobaltinitrite, 553. 

Cyanide, 330, 434. 

Dichromate, 523, 524. 

Discovery, 417. 

Electrolytic manufacture, 432. 

Ferri cyanide, 548. 

Ferro cyanide, 548. 

Fluoride, acid, 365. 

Hydroxide, 433. 

Hypochlorite, 209, 



Potassium, Iodide, 369. 

Ions, 434. 

Manganate, 520. 

Name, 433. 

Nitrate, 217, 430, 431. 

Nitrite, 430. 

Oxide, 433. 

Perchlorate, 432. 

Permanganate, 518. 

Preparation, 429. 

Properties, 429. 

Relation to life, 433, 434. 

Salts, Stassfurt, 428. 

Silicate, 383. 

Sulphate, 434. 

Sulphocyanate, 330. 

Tartrate, acid, 319, 320. 

Test, 429. 

Thiocyanate, 330. 
Pottery, 503. 
Powder, gun, 431. 

Smokeless, 328. 
Pressure, critical, 128. 

Normal, 41. 

Osmotic, 136. 

Partial, 60, 65. 
Priestley, 27, 82, 118, 173, 213, 222. 
Problems based on equations, 111. 
Producer gas, 32, 283. 
Properties, 1. 

And atomic weights, 357. 

Chemical, 5. 

Classification, 3. 

Physical, 4. 

Radioactive, 476. 
Proteid, 330. 
Protein, 218, 330. 
Proust, 90. 
Prussian blue, 549. 
Prussiate of potash, 548. 
Prussic acid, 330. 
Puddling, 538. 
Pulmotor, 26. 
Purple of Cassius, 461. 
Putty, 467. 
Pyrene, 313. 
Pyrite, 547. 
Pyroligneous acid, 317. 
Pyrolusite, 517. 
Pyromorphite, 510. 
Pyrophosphates, 397, 



588 



INDEX 



Pyrophosphoric acid, 396. 
Pyrosulphates, 351. 
Pyrosulphuric acid, 352. 
Pyrrhotite, 532. 

Quadrivalent elements, 254, 256, 257. 
Qualitative analysis, 78, 340. 
Quantitative analysis, 78, 246, 247. 
Quartation, 459. 
Quartz, 379. 
Quicklime, 468. 
Quicksilver, 487. 
Quinquivalent elements, 257. 

Radiations, 528, 529. 
Radical, denned, 154. 

Ammonium, 214. 

Organic, 311, 312. 

Valence, 255, 256. 
Radioactivity, 476, 516, 528-531. 
Radium, 476, 528-531. 

And helium, 123, 530. 

And niton, 530. 

Decay, 528. 

See Radioactivity. 
Ramsay, 118, 122, 123. 
Raoult, 240. 
Rational formula, 312. 
Rayleigh, 122. 
Rays, gamma, 529. 

X, 529. 
Reaction, defined, 101. 

Acid, 150. 

Alkaline, 151. 

Basic, 151. 

Forward, 188, 190, 192. 

Gas, 189. 

Neutral, 152. 

Reverse, 190, 191. 

Reversible, 186. 

Velocity, 184, 185, 188-191. 
Reagents, 6. 
Realgar, 401, 403. 
Red fire, 474. 

Lead, 513. 

Liquor, 319, 501. 
Reducing agent, 38. 

Carbon monoxide, 284. 

Charcoal, 271. 

Hydrogen, 38. 

Stannous chloride, 490, 509. 



Reducing flame, 308. 
Reduction, defined, 38. 

Broad meaning, 490, 509, 545. 

By charcoal, 271. 

Ionic standpoint, 545. 

Ore, 415. 
Refining petroleum, 294. 
Regenerative process, 542. 
Relative humidity, 120. 
Repulsion and attraction, 141. 
Respiration, 26. 
Reverberatory furnace, 415. 
Reverse reaction, 190, 191. 
Reversible equations, 189. 
Reversible reactions, 186-195. 
Reversion, fertilizer, 401. 
Rhigolene, 294. 
Rhodium, 554. 
Rhodocroisite, 517. 
Richards, 246. 
Rochelle salt, 320, 448: 
Roll sulphur, 334. 
Rose's metal, 407. 
Rouge, 546. 

Rubidium, 417, 435, 440. 
Ruby, 497. 
Rum, 315. 
Rusting, 544. 
Ruthenium, 554. 
Rutherford, nitrogen, 117. 

Radium, 528. 

Saccharose, 324. 
Safety lamp, 308. 
Sal ammoniac, 436. 

Soda, 422. 
Saleratus, 423. 
Salt, defined, 157, 165. 

Acid, 157, 165. 

Basic, 158, 165. 

Cake, 421. 

Common. See Sodium chloride. 

Double, 428, 501. 

Glauber's, 426. 

Microcosmic, 396. 

Normal, 157, 165. 

Pink, 509. 

Rochelle, 320, 448. 

Tartar, 433. 

Tin, 508. 
Saltpeter, Chile, 368, 369, 370, 426. 



INDEX 



589 



Salts, 152-159, 165-167, 194, 195. 

Action on litmus, 152, 165. 

And ionization, 165. 

Classification, 156. 

Composition, 153. . 

Dissociation, table, 164. 

Double, 428, 501. 

Epsom, 480. 

In sea water, 196. 

Nomenclature, 157, 161, 162. 

Preparation, 158. 

Properties, 152. 

Smelling, 438. 

Solubility product, 194, 195. 

Stassfurt, 428. 
Sand, 379, 385. 

Blast, 380. 

See Quartz. 
Saponification, 323. 
Sapphire, 497. 
Satin spar, 470. 
Saturated solution, 67. 

And equilibrium, 194, 195. 
Scheele, 25, 27, 196. 

Green, 403. 
Scrubber, 297. 
Sea water, elements in, 11. 

Solids in, 53. 
Seaweed ash, 368, 369, 370. 
Seidlitz powders, 320, 423. 
Selenite, 470. 
Selenium glass, 389. 
Semi-permeable membrane, 137. 
Serpentine, 383. 
Siderite, 532, 533, 548. 
Siemens-Martin process, 541. 
Silica, 378. 

And plants, 380. 

See Quartz. 

Vessels, 380. 
Silicates, 380, 381, 382, 383. 

Alkaline, 383, 384. 

Soluble, 379. 
Siliceous sinter, 384. 
Silicic acids, 380, 381, 382. 
Silicides, 385. 
Silicified wood, 379. 
Silicon, 378. 

Bronze, 447. 

Dioxide, 365, 378. 

Test, 384. 



Silicon, Tetrafluoride, 365, 384. 
Siloxicon, 384. 
Silver, 452. 

Alloys, 452, 453, 455. 

Bromide, 456. 

Chloride, 452, 456. 

Cleaning, 454. 

Coins, 455. 

Cyanide, 455. 

Determination atomic weight, 
247. 

German, 447. 

Horn, 196, 452. 

Iodide, 456. 

Ions, 146, 147, 456. 

Metallurgy, 453. 

Mirror, 455. 

Nitrate, 455. 

Oxidized, 454. 

Phosphates, 397. 

Plating, 455. 

Properties, 454. 

Separation from gold, 453, 459. 

Sterling, 455. 

Sulphides, 452. 

Tarnishing, 339, 454. 

Test, 456. 
Simplest formula, 249, 250. 
Slag, 415, 541. 
Smalt, 552. 
Smaltite, 552. 
Smelling salts, 438. 
Smelting, 414. See Metallurgy. 
Smithsonite, 482. 
Smokeless powder, 328. 
Soap, 321, 323. 

And hard water, 473. 

Castile, 324. 

Hard, 323. 

Manufacture, 324. 

Soft, 323. 
Soda, 422, 423. 

Ash, 422. 

Baking, 422. 

Calcined, 422. 

Caustic, 424. 

Cooking, 423. 

Crystals, 424. 

Lime, 469. 

Mints, 424. 

Washing, 422. 



590 



INDEX 



Soda water, 278. 
Sodium, 417. 

Acetate, 288, 319. 
Acid carbonate, 422. 
Acid sulphite, 344. 
Amalgam, 419, 425. 
Ammonium phosphate, 396. 
And water, 31, 61, 79. 
Arsenate, 403. 
Arsenite, 403. 
Bicarbonate, 281, 422, 423. 
Bromide, 365, 366. 
Carbonate, 417, 421, 422. 
Carbonate, hydrolysis, 166, 167, 

422. 
Carbonate, Leblanc process, 421. 
Carbonate, Solvay process, 422. 
Chloride, 419, 420, 421, 426, 428, 

504. 
Chloride, equilibrium in solution, 

193, 194. 
Chromate, 525. 
Cyanide, 330, 419, 427. 
Dichromate, 525. 
Dioxide, 427. 
Discovery, 417. 
Electrolytic manufacture, 417, 

418. 
Hydride, 419. 
Hydroxide, 424, 425. 
Hydroxide, electrolysis, 417, 

418. 
Hypochlorite, 209. 
Hyposulphite, 352. 
In crust, 11, 417. 
Iodate, 368, 370. . . 

Ions, 147, 425, 427. 
Manganate, 520. 
Manufacture, 417. 
Metaborate, 376. 
Metastannate, 507. 
Name, 417. 
Nitrate, 426, 427, 430. 
Nitrite, 427. 
Occurrence, 417. 
Organic salts, 321, 323. 
Peroxide, 418, 419, 427. 
Plumbate, 515. 
Properties, 418. 
Silicate, 379, 383. 
Btannate, 509, 



Sodium, Sulphate, 421, 426. 

Sulphate, electrolysis, 143. 

Sulphide, 421. 

Sulphite, 426. 

Sulphite, acid, 344, 370. 

Test, 418. 

Thiosulphate, 352, 457. 

Tungstate, 527. 

Uranate, 527. 

Uses, 419. 
Soft coal, 269. 

Water, 473. 
Solder, 508. 
Soldering, 376. 
Solubility, defined, 67. 
■ Curve, 68. 

Degree, 63. 

Product, 195. 
Solute, 63. 
Solution, 63-76, 132-148, 193-195. 

And chemical action, 76, 184. 

And crystallization, 70. 

And dissociation, 163. 

And electric current, 132. 

And electrolysis, 140. 

And equilibrium, 189, 193-195. 

And osmotic pressure, 136. 

And vapor pressure, 73. 

Boiling point, 140. 

Colloids, 382. 

Concentration and reaction, 185, 
193-195. 

Dissociation table, 164. 

Electrolysis, 140. 

Electrolytic, 133. 

Freezing point, 139, 239, 240. 

Gases, 64. 

General properties, 132. 

Heat of, 75. 

Liquids, 65. 

Non-electrolytic, 133. 

Normal, 163. 

Pressure, 452. 

Saturated, 67, 194, 195. 

Solids, 66-76. 

Summary, 147. 

Supersaturated, 69. 

Tension, 451. 

Terms, 63, 64. 

Thermal phenomena, 75, 
Solvay process, 422, 



INDEX 



591 



Solvent, 63. 

Spathic iron ore, 548. 

Specific gravity, metals, 412. 

Water, 56. 
Specific heat, 245. 

And atomic weight, 246. 

Law, 245. 

Table, 246. 
Spectra, inert gases, 124. 
Spectroscope, 439. 
Spectrum, 439. 

Analysis, 438. 

Barium, 475. 

Calcium, 472. 

Line, 440. 

Potassium, 440. 

Radium, 476. 

Sodium, 440. 

Strontium, 474. 
Speculum metal, 447. 
Spelter, 482. 
Sperrylite, 554. 
Sphalerite, 482. 
Spiegel iron, 517. 
Spinels, 497. 
Spinthariscope, 529. 
Spontaneous combustion, 24. 
Stahl, 25. 
Stalactite, 465. 
Stalagmite, 465. 
Standard conditions, 42, 47. 
Stannic compounds, 508, 509. 
Stannous compounds, 490, 508, 509. 
Starch, 327. 

And alcohol, 315. 

And iodine, 327. 

Test, 370. 
Stas, 357. 
Stassfurt deposits, 428. 

Bromine, 366. 

Calcium chloride, 472. 

Chlorine, 196. 

Magnesium, 478. 

Minerals, 428. 

Potassium, 428. 

Sodium sulphate, 426. 
Steam, 58. 
Stearin, 321, 322. 
Steel, 538. 

Acid process, 541. 

Aluminium, 495. 



Steel, Basic process, 540, 541. 

Bessemer process, 539. 

Cementation process, 542. 

Composition, 538. 

Crucible process, 542. 

Important properties, 543. 

Open hearth, 541. 

Siemens-Martin process, 541. 

Special, 543. 

Tempering, 543. 

Thomas-Gilchrist process, 540. 

Uses, 543. 
Sterling silver, 455. 
Stibine, 405. 
Stibnite, 405, 406. 
Stoneware, 504. 
Storage cell, 180. 
Stove polish, 265. 
Strass, 388. 
Strontia, 474. 
Strontium, 473, 474, 475. 
Structural formulas, 260, 312. 
Stucco, 471. 
Sublimate, 436. 

Corrosive, 489. 
Sublimation, ammonium chloride, 
436. 

Iodine, 370. 
Substitution, 32, 61, 103. 
Sucrose, 324. 
Sugar, 324. 

And osmotic pressure, 137. 

Beet, 325. 

Cane, 324. 

Fermentation, 315, 326. 

Fruit, 326. 

Granulated, 325. 

Grape, 326. 

Manufacture, 325. 

Molecular weight, 241. 

Of lead, 319. 

Refining, 325. 

Solution, freezing point, 139. 

Test, 326. 
Suint, 428. 

Sulphides, 332, 339, 340, 399, 437. 
Sulphites, 344. 

Ions, 344, 345. 
Sulphur, 332. 

Amorphous, 336, 

Chlorides, 335. 



592 



INDEX 



Sulphur, Crystallized, 336. 

Dioxide, 340, 341, 342, 426. 

Extraction, 333. 

Flowers, 334. 

Forms, 336. 

Hexoxide, 343. 

Kiln, 333. 

Lime, 337. 

Milk of, 337. 

Molecule, 335. 

Monoclinic, 336. 

Orthorhombic, 336. 

Phosphorus, 399. 

Properties, 2, 334. 

Purification, 334. 

Roll, 334. 

Spray, 337. 

Transition point, 336. 

Trioxide, 343. 

Winning, 333. 
Sulphuretted hydrogen, 337. 
Sulphuric acid, 345. 

Calculation formula, 100. 

Chamber process, 346. 

Contact process, 348. 

Di-, 352. 

Fuming, 351. 

Ions, 350, 351. 

Manufacture, 345. 

Nordhausen, 351. 

Plant, 346, 347, 348. 

Properties, 349. 

Pyro-, 352. 

Test, 351. 

Uses, 350. 
Sulphuric ether, 316. 
Sulphurous acid, 342. 

Ions, 344. 
Sun, atmosphere, 30. 
Superphosphate of lime, 400. 
Supersaturation, 69. 
Sylvite, 428. 
Symbols, 10, 13, 97. 

And ions, 134. 

Latin, 11. 

Tables, 10, 11, 12, 564. 
Synthesis, water, 81. 

Table salt, 419. 

Tables, atomic weights, 231, 564. 
Borax beads, 376. 



Tables, Calorific value of coal, 269. 

Combination by volume, 225. 

Combining water, 71. 

Composition of air, 118. 

Composition of body, 12. 

Composition of earth's crust, 11. 

Composition of illuminating gas, 
300. 

Composition of ocean, 11, 196. 

Dissociation, 164. 

Distribution of matter, 11. 

Equivalent weights, 230, 231, 
232. 

Important elements, 10. 

International atomic weights, 
564. 

Ionization, 164. 

Ions, 147. 

Latin symbols, 13. 

Less common elements, 12. 

Metalloids, 411. 

Metals, 411. 

Metric equivalents, 559. 

Metric system, 559. 

Metric transformations, 559. 

Multiple proportions, 91. 

Nitrogen oxides, 221. 

Non-metals, 411. 

Periodic, 358. 

Solids in ocean, 53. 

Solubility of solids, 67. 

Specific heat, 246. 

Valence, elements, 254, 256. 

Valence, radicals, 256. 

Vapor pressure, 563. 

Water in food, 50. 
Tar, coal, 295, 297, 329. 
Tartar, 319. 

Emetic, 320, 406. 
Tartaric acid, 319. 
Tellurides, 458. 
Temperature, critical, 128. 

Kindling, 25, 172. 

Low, ethylene, 289. 

Standard, 41. 
Tempering, 543. 
Temporary hardness, 473. 
Tension, water vapor, 58. 
Test, 21, 145, 208. 

Acetic acid, 320. 

Alcohol, 320. 



INDEX 



593 



Test, Aluminium, 502. 

Antimony, 406. 

Arsenic, 404. 

Barium, 475. 

Bismuth, 407. 

Borax bead, 376. 

Boron, 375. 

Cadmium, 486. 

Calcium, 472. 

Carbon, 272. 

Carbon dioxide, 276, 470. 
CChloride, 145. 

Chlorine ions, 145. 

Chromium, 524, 525. 

Cobalt, 553. 

Copper, 450. 

Gold, 461. 
.Hydrogen, 37. 

Hydrogen sulphide, 340. 

Iodine, 327, 370. 

Iron, 549. 

Lead, 514, 515. 

Lithium, 435. 

Magnesium, 481. 

Manganese, 518, 520. 

Marsh's, for arsenic, 404. 

Mercury, 490. 

Nickel, 551. 

Nitrates, 220. 

Nitric acid, 220. 

Oxygen, 21. 

Potassium, 429. 

Silicon, 384. 

Silver, 456. 

Sodium, 418. 

Starch, 327, 370. 

Strontium, 474. 

Sugar, 326. 

Sulphate ions, 145. 

Sulphates, 351. 

Sulphuric acid, 351. 

Tin, 509. 

Zinc, 485. 
Tetraboric acid, 375. 
Tetrads, 257. 
Thenard, 87. 
Theory, 88. 

Atomic, 91. 

Electrolytic dissociation, 
180. 
Thermal equation, 176, 275. 



133, 



Thermal, Phenomena of solution, 75. 
Thermit, 496. 
Thermometer, 560. 
Thomas-Gilchrist process, 540. 

Slag, 541. 
Thorium, 515, 516, 531. 
Tin, 505. 

Alloys, 508. 

Amalgam, 508. 

Chlorides, 508, 509. 

Compounds, 508. 

Crystals, 508. 

Dioxide, 505, 506, 508. 

Disease, 506, 507. 

Displacement, 507. 

Foil, 508. 

Ions, 509. 

Metal and non-metal, 509. 

Mordants, 509. 

Oxymuriate, 509. 

Properties, 506. 

Salt of, 508. 

Stone, 505. 

Stream, 508. 

Sulphides, 509. 

Test, 509. 

Transition point, 506. 
Topaz, 497. 
Transition point, 336. 

Sulphur, 336. 

Tin, 506. 
Transmutation, 458. 
Travertine, 465. 
Triacid base, 158. 
Triads, 257. 
Tribasic acid, 158. 
Tricalcium phosphate, 391-394, 396, 

400. 
Trisilicate, 383. 
Trivalent elements, 257. 
Tungsten, 527. 
Turnbull's blue, 549. 
Turquoise, 497. 
Tuyeres, 534. 
Type metal, 512. 

Univalent elements, 254, 256, 257. 
Upward displacement, 211. 
Uranium, 476, 527-531. 

Products, 531. 
Urea, 310. 



594 



INDEX 



Valence, 252, 253. 

And atomic weight, 254. 

And equivalent weight, 254. 

Determination, 253. 

Exceptions, 260. 

Radicals, 255. 

Raising and lowering, 545. 

Representation, 259, 312. 

Rules, 257, 258, 259. 

Tables, 254, 256. 

Variable, 256. 
Valentine, Basil, 345. 
Vanadium, 543. 
Van Helmont, 282. 
Vanillin, 316. 
Vapor density, 236. 

And molecular weight, 236-238. 

Determination, 238, 239. 

Method, 236, 239. 
Vapor pressure, 58-61. 

And deliquescence, 74, 424. 

And efflorescence, 72, 422. 

And solution, 73. 

Table, 563. 
Vapors, molecular formula, 250. 
Vapor tension, 58. 
Vaseline, 294. 

Velocity, reaction, 184, 185, 188-191. 
Venetian red, 546. 
Verdigris, 319, 450. 
Vermilion, 489. 
Victor Meyer, 239. 
Vinegar, 317, 318. 
Vital force, 310. 
Vitriol, blue, 448. 

Green, 546. 

Oil of, 345. 

White, 484. 
Volatile alkali, 161. 
Volta, 176. 
Voltaic cell, 178. 
Volume equation, 252. 
Volumetric composition, 81. 

Air, 118. 

Water, 81. 

Washing soda, 422. 
Water, 50-85. 

Analysis, 54. 

And calcium oxide, 62. 

And electric current, 62, 143. 



Water, And ether, 66. 

And hydrogen, 31, 36, 78. 

And neutralization, 155. 

And oxygen, 79. 

And potassium, 429. 

And sodium, 31, 61, 79. 
. -Baryta, 474. 

Bromine, 367. 

Catalysis, 186, 200. 

Chalybeate, 53. 

Chemical properties, 61. 

Chlorine, 198. 

Composition, 78-85. 

Crystallization, 71, 98. 

Decomposition, 61, 80, 419. 

Density, 56. 

Dissociation, 165, 194, 422, 448, 
449. 

Distillation, 54. 

Drinking, 53. 

Electrolysis, 143. 

Evaporation, 57-61. 

Expansion, 56. 

Gas, 283, 298, 299, 300. 

General properties, 50-76. 

Glass, 383. 

Gravimetric composition, 83. 

Hard, 281. 

Hardness, 473. 

Hydrogen sulphide, 338. 

Hydrolysis, 166. 

Industrial application, 51. 

In food, 50. 

In nature, 50, 51. 

Ions, 165, 194, 422, 448, 449. 

Javelle's, 209. 

Lime, 469. 

Lithia, 435. 

Mineral, 52. 

Natural, 52. 

Ocean, 11, 53. 

Of crystallization, 71, 98. 

Pure, properties, 56. 

Purification, 499. 

River, 53. 

Royal, 224. 

Soda, 278. 

Soft, 473. 

Solvent power, 63. 

Underground, 52, 281, 465. 

Vapor and catalysis, 86. 



INDEX 



595 



Water, Vapor in atmosphere, 120. 

Vapor, molecular equation, 251 
252. 

Vapor. See Vapor pressure. 

Volumetric composition, 81. 
Watt, 83. 
Weight, liter of air, 115. 

Ammonia gas, 211. 

Carbon dioxide, 277. 

Chlorine, 198. 

Formula, 238. 

Hydrogen, 33. 

Hydrogen chloride, 205. 

Hydrogen sulphide, 338. 

Molar, 238. 

Oxygen, 19, 48, 59. 

Nitrogen, 129. 

Sulphur dioxide, 342. 
Weldon process, 197, 518. 
Welsbach light, 309, 516. 
Whisky, 315. 
White arsenic, 402. 

Lead, 513, 514. 

Metal, 508. 

Paint, 513. 

Vitriol, 484. 
Whitewash, 470. 
Whiting, 467. 
Willemite, 476, 482. 
Willson, 285. 
Wine, 315. 
Witherite, 474. 
Wohler, 310. 
Wollastonite, 383. 
Wood alcohol, 314. 
Ashes, 428, 432. 



Wood, Charcoal, 271, 272, 273. 

Distillation, 272, 317. 

Petrified, 379. 

Silicified, 379. 

Vinegar, 317. 
Wood's metal, 486. 
Wrought iron, 537. 

X-rays, 529. 
Xenon, 123, 358. 

Yeast, 315, 327. 

Zero, absolute, 44, 560. 
Zinc, 482. 

Alloys, 484. 

Blende, 482. 

Chloride, 484. 

Determination atomic weight, 
247. 

Displacement, 451, 483. 

Dust, 482, 484. 

Family, 491. 

Hydroxide, 485. 

Ions, 451, 485. 

Oxide, 482, 483, 484. 

Properties, 482, 485. 

Silicate, 482. 

Sulphate, 451, 484. 

Sulphide, 476, 482, 484, 529. 

Test, 485. 

White, 484. 
Zincates, 483, 485. 
Zincite, 482. 
Zircon, 383. 
Zymase, 315. 



I 



International and Approximate Atomic Weights 



• 

Element 


Symbol 


Inter. 


Approx. 

At. Wt, 


Aluminum ........ 


Al 


27.1 


27 


Antimony 


Sb 


120.2 


120 


Arsenic 


As 


74.96 


75 


Barium 


Ba 


137.37 


137 


Bismuth * 


Bi 
B 


208.0 
11.0 


208 


Boron 


11 


Bromine 


Br 


79.92 


80 


Cadmium 


Cd 


112.40 


112 


Calcium 


Ca 


40.07 


40 


Carbon 


C 


12.005 


12 


Chlorine 


CI 


35.46 


35.5 


Chromium 


Cr 


52.0 


52 


Cobalt 


Co 


58.97 


59 


Copper 


Cu 


63.57 


63.5 


Fluorine • 


F 


19.0 


*19 


Gold 


Au 


197.2 


197 


Hydrogen 


H 


1.008 


1 




I 


126.92 


127 


Iron 


Fe 


55.84 


56 




Pb 


207.20 


207 


Lithium 


Li 


6.94 


7 




Mg 


24.32 


24 


Manganese 


Mn 


54.93 


55 




Hg 


200.6 


200.5 


Nickel 


Ni 


58.68 


58.5 




N 


14.01 


14 


Oxygen 





16.00 


16 


Phosphorus 


P 


31.04 


31 


Platinum 


Pt 


195.2 


195 




K 


39.10 


39 


Radium 


Ra 

Si 


226.0 

28.3 


226.0 




28 


Silver 


Ag 


107.88 


108 




Na 


23.00 


23 


Strontium 


Sr 


87.63 


87.5 


Sulphur 


S 


32.06 


32 


Tin 


Sn 
Zn 


118.7 
65.37 


118.5 




65 



The international atomic weights, given above, are taken from the 
1916 table. The approximate atomic weights are supplied by the 
author and are designed to facilitate the solution of problems. 



LIBRARY OF CONGRESS 





003 610 298 6 



